Edexcel Chemistry AS Notes
Edexcel Chemistry AS Notes
Edexcel Chemistry AS Notes
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1 unit of charge is 1.602 x 10-19 coulombs. A proton is given a charge of +1 and an electron a charge of -1. All charges are measured in these units. 1 unit of mass is 1.661 x 10-27 kg. This is also not a convenient number, so we use atomic mass units. Since the mass of protons and neutrons varies slightly depending on the nucleus, then in order to define an atomic mass unit we need to choose one nucleus as a standard. For this purpose 126C , or carbon-12, was chosen because its mass per nucleon (1.661 x 10 27 kg) is around average, which means all the other nuclei have masses close to whole numbers. An atomic mass unit is thus defined as 1/12th of the mass of one atom of carbon-12. Everything else is measured relative to this quantity.
b) Atomic numbers, mass numbers and isotopes An atom is named after the number of protons in its nucleus. If the nucleus of an atom has 1 proton, it is hydrogen; if it has two protons, it is helium; if it has 3, it is lithium etc. The number of protons in the nucleus of an atom is called the atomic number. It has the symbol Z. The atomic number is the number of protons in the nucleus of an atom
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3 Not all atoms of the same element have equal numbers of neutrons; this may vary slightly. The sum of the number of protons and neutrons in the nucleus of an atom is called its mass number. It is represented by the symbol A. The mass number is the sum of the number of protons and neutrons in the nucleus of an atom
The nucleus of an atom can thus be completely described by its mass number and its atomic number. It is generally represented as follows:
A ZE
Eg. 94Be, 126C, 2412Mg Atoms with the same atomic number but with different mass numbers (ie different numbers of neutrons) are called isotopes. Isotopes are atoms with the same atomic number but with different mass numbers Eg magnesium (atomic number 12) has 3 naturally occurring isotopes:
24 12Mg: 25 12Mg: 26 12Mg:
In a neutral atom, the number of protons and electrons are the same. However, many elements do not exist as neutral atoms, but exist as ions. Ions are species in which the proton and electron numbers are not the same, and hence have an overall positive or negative charge. The number of electrons in a species can be deduced from its charge: Eg
24 2+ 12Mg : 12p, 12n, 10e 24 + 12Mg : 12p, 12n, 11e 24 12Mg 12p, 12n, 12e 24 12Mg : 12p, 12n, 13e
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4 c) Relative atomic mass The mass of an atom is measured in atomic mass units, where one unit is 12th of the mass of one atom of carbon-12. The relative isotopic mass of an isotope is the ratio of the mass of one atom of that isotope to 1/12th of the mass of one atom of carbon-12. It is usually very close to a whole number ratio: Isotope
1 1H 4 2He 9 4Be 27 13Al 59 27Co
Mass number 1 4 9 27 59
The masses of protons and neutrons vary slightly from isotope to isotope, so the relative isotopic mass is not exactly a whole number. The relative atomic mass of an atom is the ratio of the average mass of one atom of that element to 1/12th of the mass of one atom of carbon-12. The RAM is the average mass of all the isotopes, and is often not close to a whole number: Element Mg Cl Br Ba Common mass numbers 24, 25, 26 35, 37 79, 81 134, 135, 136, 137, 138 Relative atomic mass 24.32 35.45 79.91 137.33
Some elements and compounds exist as molecules; these also have a characteristic mass: The relative molecular mass of a molecule is the ratio of the average mass of that molecule to 1/12th of the mass of an atom of carbon-12. The relative molecular mass of a molecule is the sum of the relative atomic masses of its constituent atoms. Eg The relative molecular mass of CO2 is 12.0 + 16.0 + 16.0 = 44.0
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5 MASS SPECTROMETRY The mass spectrometer is an instrument used for measuring the masses of atoms and molecules. It can also be used to measure the relative abundance of different isotopes and to predict the structure of more complex molecules. 1. How the mass spectrometer works
The workings of the mass spectrometer can be summarized in five stages: 1- Gaseous material released into ionization chamber 2- Particles bombarded with electrons and ionized, mostly to +1 ions (IONISATION) A metal wire is heated until it starts emitting high energy electrons. These electrons hit the particles, knocking more electrons off. Most of the particles are ionized to +1 ions 3- Ions accelerated to uniform speed by electric field (ACCELERATION) The positive ions are attracted to the negative plate and accelerate towards it 4- Ions deflected by magnetic field; deflection (DEFLECTION) The heavier the particle, the less the deflection depends on m/e ratio
5- Electric current measured as ions land on plate (DETECTION) The greater the abundance of the isotope, the larger the current
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6 The degree of deflection depends on the mass and the charge; the greater the mass, the less the deflection, and the greater the charge, the greater the deflection. It can be shown that the deflection is inversely proportional to the m/e ratio. In most cases, however, the charge is +1, so the deflection depends essentially on the relative mass of the species in the mass spectrometer. If the spectrometer is calibrated, the masses of all the species can be directly measured. The greater the number of particles landing at a single point on the detector, the greater the electric current and the larger the peak. Thus the relative abundance of different isotopes can be measured. Since the position at which an ion appears on the detector depends on its mass, different isotopes appear at different points on the detector. The magnitude of the peak gives the relative abundance of the isotope. Thus the relative atomic mass of the element can be calculated from its mass spectrum. An example of a simple mass spectrum is shown below: Mass spectrum of Ne
100
80
relative abundance
60
40
20
18
20
22 M/Z
24
26
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7 2. Calculating relative atomic masses The relative atomic mass can be calculated by the formula: (perentage abundance of each isotope x mass of each isotope) 100 Eg. Using the mass spectrum of neon above: RAM = (90 x 20 + 10 x 22)/100 = 20.2 All relative atomic masses have been found in this way.
3. Deducing relative molecular masses It is also possible to put molecules into the mass spectrometer. Because the conditions inside a mass spectrometer are very extreme, the molecules often break up into smaller pieces. This is known as fragmentation. The mass spectrum of a molecule can thus look quite complicated: Mass spectrum of pentane (C5H12)
Many of these peaks result from fragmentation of the molecule, but the peak with the largest m/e ratio comes from the unbroken molecular ion, in this case C 5H12+, and is called the molecular ion peak. The m/e ratio of this peak (72) will be the relative molecular mass of the molecule. The relative molecular mass of a molecule is obtained by looking at the peak in the spectrum with the largest m/e ratio (ie the peak furthest to the right).
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8 ELECTRONIC STRUCTURE i) Energy levels Electrons do not orbit the nucleus randomly; they occupy certain fixed energy levels. Each atom has its own unique set of energy levels, which are difficult to calculate but which depend on the number of protons and electrons in the atom. Energy levels in an atom can be numbered 1,2,3,. To infinity. 1 is the lowest energy level (closest to the nucleus) and energy level infinity corresponds to the energy of an electron which is not attracted to the nucleus at all. The energy levels thus converge as they approach infinity:
n=2
n=1
ii) Orbitals and sub-levels Electrons do not in fact orbit the nucleus in an orderly way. In fact they occupy areas of space known as orbitals. The exact position of an electron within an orbital is impossible to imagine; an orbital is simply an area of space in which there is a high probability of finding an electron. Orbitals can have a number of different shapes, the most common of which are as follows:
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Every energy level contains one s-orbital. An s-orbital in the first energy level is a 1s orbital. An s-orbital in the second energy level is a 2s orbital, etc
p-orbitals: these are shaped like a 3D figure of eight. They exist in groups of three:
Every energy level except the first level contains three p-orbitals. Each p-orbital in the same energy level has the same energy but different orientations: x, y and z. A p-orbital in the second energy level is a 2p orbital (2px, 2py, 2pz) A p-orbital in the third energy level is a 3p orbital (3px, 3py, 3pz), etc In addition, the third and subsequent energy levels each contain five d-orbitals, the fourth and subsequent energy levels contain seven f-orbitals and so on. Each type of orbital has its own characteristic shape. S, p and d orbitals do not all have the same energy. In any given energy level, sorbitals have the lowest energy and the energy of the other orbitals increases in the order p < d < f etc. Thus each energy level must be divided into a number of different sub-levels, each of which has a slightly different energy. The number and type of orbitals in each energy level can thus be summarised as follows: Energy level Number and type of orbital 1st sublevel 1 x 1s 1 x 2s 1 x 3s 1 x 4s 1 x 5s 2nd sub- 3rd sub- 4th sub- 5th sublevel level level level 3 x 2p 3 x 3p 3 x 4p 3 x 5p
1 2 3 4 5
5 x 3d 5 x 4d 5 x 5d
7 x 4f 7 x 5f
9 x 5g
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10 iii) Shells Since the different sub-levels have different energies, and the energies of the different levels get closer together with increasing energy level number, the high energy sub-levels of some energy levels soon overlap with the low energy sublevels of higher energy levels, resulting in a more complex energy level diagram:
n=4 n=3 E N E R G Y
4f 4d 4p 3d 4s 3p 3s 2p 2s
n=2
n=1
1s
Starting with the lowest energy, the orbitals can thus be arranged as follows: 1s 2s 2p 3s 3p 5d 6p 7s 4s 3d 5f 6d 4p 5s 4d 5p 6s 4f
Many of these sub-levels have similar energy, and can be grouped together. A collection of sub-levels of similar energy is called a shell. 1s2s 2p3s 3p 4s 3d 4p 5s 4d 5p6s 4f 5d 6p
The arrangement of shells and the maximum number of electrons in each can be summarised as follows: Shell number 1 2 3 4 5 6 Orbitals in shell 1 x1s 1 x 2s, 3 x 2p 1 x 3s, 3 x 3p 1 x 4s, 5 x 3d, 3 x 4p 1 x 5s, 5 x 4d, 3 x 5p 1 x 6s, 7 x 4f, 5 x 5d, 3 x 6p
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11 iv) Electrons Electrons repel each other. In a small space such as an orbital, it is impossible to put more than two electrons. Since electrons are charged particles, and moving charges create a magnetic field, it is possible to create a small magnetic attraction between two electrons if they are spinning in opposite directions in the same orbital. This is the reason two electrons, and not one, are permitted in the same orbital. It is thus possible to calculate the maximum possible number of electrons in each sub-level, and thus in each energy level: Shell 1 2 3 4 5 6 Number of electrons in each sub-level 2 x 1s 2 x 2s, 6 x 2p 2 x 3s, 6 x 3p 2 x 4s, 10 x 3d, 6 x 4p 2 x 5s, 10 x 4d, 6 x 5p 2 x 6s, 14 x 4f, 10 x 5d, 6 x 6p Max. no of electrons 2 8 8 18 18 32
v) Electron arrangement in orbitals There are three rules which determine the way in which electrons fill the orbitals 1. Aufbau/building principle: electrons always fill the lowest energy orbitals first. 2. Hund's rule: electrons never pair up in the same orbital until all orbitals of the same energy are singly occupied, and all unpaired electrons have parallel spin. 3. Pauli exclusion principle: only two electrons may occupy the same orbital, and they must do so with opposite spin. The arrangement of electrons in an atom is known as its electronic configuration. It can be represented in two ways: The arrow and box method represents each orbital as a box and each electron as an arrow. The direction of spin is shown by the orientation of the arrow.
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12 The electronic configuration of the first 18 elements using the arrow in box method is as follows: 1s 2s 2p 3s 3p
H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar
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13 The orbital method indicates the number of electrons in each orbital with a superscript written immediately after the orbital. The electronic configurations of the first eighteen elements can be shown with the orbital method as follows: H: He: Li: Be: B: C: N: O: F: Ne: Na: Mg: Al: Si: P: S: Cl: Ar: 1s1 1s2 1s22s1 1s22s2 1s22s22p1 1s22s22p2 or 1s22s22p63s23px13py1 1s22s22p3 or 1s22s22p63s23px13py13pz1 1s22s22p4 or 1s22s22p63s23p23px23py13pz1 1s22s22p5 1s22s22p6 1s22s22p63s1 1s22s22p63s2 1s22s22p63s23p1 1s22s22p63s23p2 or 1s22s22p63s23px13py1 1s22s22p63s23p3 or 1s22s22p63s23px13py13pz1 1s22s22p63s23p4 or 1s22s22p63s23px23py13pz1 1s22s22p63s23p5 1s22s22p63s23p6
A shorthand form is often used for both the above methods. Full shells are not written in full but represented by the symbol of the element to which they correspond, written in square brackets. Eg. 1s22s22p6 is represented as [Ne] and 1s22s22p63s23p6 is represented as [Ar]. The shorthand electronic configuration of the elements with atomic numbers 18 to 36 can be written as follows:
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14 4s 3d 4p
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
[Ar] [Ar] [Ar] [Ar] [Ar] [Ar] [Ar] [Ar] [Ar] [Ar] [Ar] [Ar] [Ar] [Ar] [Ar] [Ar] [Ar] [Ar]
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15 Note the unusual structures of chromium and copper. The difference in energy between the 3d and 4s electrons is very small, and in chromium the energy required to promote and electron from 4s to 3d is recovered in the reduced repulsion which results from the fact that they are no longer paired. Thus the 4s13d5 structure in Cr is preferred. In copper, the 3d orbitals are actually lower in energy than the 4s orbital, so the 4s13d10 structure in Cu is preferred. v) Electron arrangement in ions The electronic configuration of ions can be deduced by simply adding or removing the appropriate number of electrons. The order in which electrons are to be removed can be deduced from the following rules: remove outer shell electrons first remove p-electrons first, then s-electrons and then d-electrons remove paired electrons before unpaired electrons in the same sub-level
vi) Effect of electronic configuration on chemical properties The chemical properties of an atom depend on the strength of the attraction between the outer electrons and the nucleus. These in turn depend on the number of protons and on the electronic configuration, and so it follows that these two factors are instrumental in determining the chemical properties of an atom. This is in contrast with the neutron number however, which has no effect on the chemical properties of an atom. Neutrons have no charge and hence exert no attractive force on the nucleus. Isotopes, therefore, tend to have very similar chemical properties since they have the same atomic number and the same electronic configuration. They differ only in number of neutrons, which do not directly influence the chemical properties of an element.
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16
IONISATION ENERGIES
i) First ionisation energy The first ionisation energy of an element is the energy required to remove one electron from each of a mole of free gaseous atoms of that element. It can also be described as the energy change per mole for the process: M(g) M+(g) + e The amount of energy required to remove an electron from an atom depends on the number of protons in the nucleus of the atom and on the electronic configuration of that atom. The first ionisation energies of the first 20 elements in the periodic table is shown below:
Variation of first ionisation energy with atomic number for the first twenty elements
first ionisation energy (kJ per mole)
2500 2000 1500 1000 500 0 0 5 10 atomic number 15 20
There are various trends in this graph which can be explained by reference to the proton number and electronic configuration of the various elements. A number of factors must be considered: - Energy is required to remove electrons from atoms in order to overcome their attraction to the nucleus. The greater the number of protons, the greater the attraction of the electrons to the nucleus and the harder it is to remove the electrons. The number of protons in the nucleus is known as the nuclear charge. - The effect of this nuclear charge, however, is cancelled out to some extent by the other electrons in the atom. Each inner shell and inner sub-shell electron effectively cancels out one unit of charge from the nucleus. This is known as shielding.
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17 - The outermost electrons in the atom thus only feel the residual positive charge after all inner shell and inner sub-shell electrons have cancelled out much of the nuclear charge. This residual positive charge is known as the effective nuclear charge. - Electrons repel each other, particularly when they are in the same orbital. The degree of repulsion between the outermost electrons affects the ease with which electrons can be moved. When considering trends in ionisation energies, it is thus necessary to consider 4 factors: nuclear charge shielding effective nuclear charge electron repulsion
The trends in first ionisation energies amongst elements in the periodic table can be explained on the basis of variations in one of the four above factors.
Trend across period 1 Compare the first ionisation energies of H and He. Neither have inner shells, so there is no shielding. He has two protons in the nucleus; H only has one. Therefore the helium electrons are more strongly attracted to the nucleus and hence more difficult to remove. The first ionisation energy of He is thus higher than that of H. Since H and He are the only atoms whose outer electrons are not shielded from the nucleus, it follows that He has the highest first ionisation energy of all the elements. All elements (except H) have outer electrons which are shielded to some extent from the nucleus and thus are easier to remove. So Helium has the highest first ionisation energy of all the elements. Trends across period 2 Compare now the first ionisation energies of He (1s2) and Li (1s22s1). Li has an extra proton in the nucleus (3) but two inner-shell electrons. These inner-shell electrons cancel out the charge of two of the protons, reducing the effective nuclear charge on the 2s electron to +1. This is lower than the effective nuclear charge on the He 1s electrons, +2, and so the electrons are less strongly held and easier to remove. The first ionisation energy of Li is thus lower than that of He.
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18 Compare the first ionisation energies of Li (1s22s1) and Be (1s22s2). Be has one more proton in the nucleus than Li, and no extra inner-shell electrons, so the effective nuclear charge on Be is higher and the Be electrons are more strongly attracted to the nucleus. The first ionisation energy of Be is thus higher than that of Li. In general, the first ionisation energy increases across a period because the nuclear charge increases but the shielding remains the same. Compare the first ionisation energies of Be (1s22s2) and B (1s22s22p1).B has one more proton in the nucleus than Be but there are also 2 extra inner sub-shell electrons. These cancel out the charge of two more of the protons, leaving an effective nuclear charge of only +1. This is less than Be (+2) so the electrons are less strongly attracted to the nucleus and thus less difficult to remove. The first ionisation energy of B is thus lower than that of Be. Ionisation energies decrease from group II to group III because in group III the electrons are removed from a p-orbital, so it is shielded by the s-electrons in the outer shell. Thus the effective nuclear charge decreases. From B (1s22s22p1) to N (1s22s22p3) the proton number increases, but the number of electrons shielding the nuclear charge remains the same at 4. Thus the effective nuclear charge increases from B to N and the electrons become progressively harder to remove. The first ionisation energy thus increases from B to N.
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19 So far the concepts of effective nuclear charge and shielding have been used to explain the trend in first ionisation energies for the first 7 elements. They cannot, however, explain the fall between N and O. The electronic configurations of N and O must be considered more carefully: 1s 2s 2p N O
Note that in N the electron is removed from an unpaired orbital, but in O it is removed from a paired orbital. In a paired orbital, the two electrons share a confined space and so repel each other. They are therefore less stable and easier to remove. This repulsion effect outweighs the higher effective nuclear charge in O. The first ionisation energy of O is thus lower than that of N. First ionisation energies decrease from group V to group VI, since the electron removed from the group VI atom is paired, so there is more repulsion between the electrons and the electron is easier to remove. The first ionisation energies increase as expected from O to Ne, due to the increase in effective nuclear charge.
The trend in first ionisation energies across period 2 can thus be summarised as follows: 1. There is a general increase across the period as the nuclear charge increases and the shielding remains the same. 2. There is a drop from Be to B because in B a 2p electron is being removed and the extra shielding from the 2s subshell actually causes a fall in the effective nuclear charge. 3. There is also a drop from N to O because the electron in O is being removed from a paired orbital. The repulsion of the electrons in this orbital makes them less stable and easier to remove. The same trend can also be found in Period 3 (Na - Ar). There is a general increase, but a drop between Mg and Al and also between P and S.
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20 Trend down a group The above graph also shows a clear decrease in first ionisation energy on descending a group. This can be explained in the following way: On descending a group, the effective nuclear charge stays the same but the number of inner shells increases. The repulsion between these inner shells and the outer electrons makes them less stable, pushes them further from the nucleus and makes them easier to remove.
ii)
The second ionisation energy of an atom is the energy required to remove one electron from each of a mole of free gaseous unipositive ions. M+(g) M2+(g) + e Other ionisation energies can be defined in the same way: The third ionisation energy of an atom is the energy required to remove one electron from each of a mole of bipositive ions. M2+(g) M3+(g) + e The nth ionisation energy can be defined as the energy required for the process M(n-1)+(g) Mn+(g) + e It always becomes progressively more difficult to remove successive electrons from an atom; the second ionisation energy is always greater than the first, the third always greater than the second and so on. There are two satisfactory explanations for this: As more electrons are removed from an atom, the number of electrons remaining in the atom decreases. The repulsion between these electrons therefore decreases, while the number of protons remains the same. The remaining electrons are thus more stable and increasingly difficult to remove. The difference in successive ionisation energies, however, varies widely and depends on the electronic configuration of the atom in question. The difference in successive ionisation energies of an atom can be predicted qualitatively by consideration of the effective nuclear charge on the electron to be removed and the shielding of that electron by the inner shell and inner sub-shell electrons.
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21 Consider the successive ionisation energies of aluminium, 1s22s22p63s23p1: The 1st ionisation energy is fairly low because the 3p electron is shielded by all the other electrons, and the effective nuclear charge is only +1. The 2nd and 3rd ionisation energies are significantly higher than the 1st because 3s electrons are being removed and the effective nuclear charge on these electrons is +3. 1st: 578 kJmol-1, 2nd: 1817 kJmol-1, 3rd: 2745 kJmol-1 There is a huge jump to the 4th ionisation energy, since a 2p electron is now being removed. The shielding has fallen and the effective nuclear charge has risen to +9. The 5th and 6th ionisation energies are also high. 4th: 11578 kJmol-1, 5th: 14831 kJmol-1, 6th: 18378 kJmol-1 There is another significant jump to the 7th ionisation energy, since an unpaired 2p electron is now being removed. 7th: 23296 kJmol-1, 8th: 27460 kJmol-1, 9th: 31862 kJmol-1 The next significant jump is between the 9th and 10th ionisation energies, since the 10th requires the removal of a 2s electron. 10th: 38458kJmol-1, 11th: 42655 kJmol-1 There is a huge jump to the12th ionisation energy, since a 1s electron is now being removed. 12th: 201276kJmol-1, 13th: 222313kJmol-1. These ionisation energies could be plotted on a graph as follows:
ionisation energy
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22 Note that the largest jumps by far occur between the 3rd and 4 th ionisation energies, and between the 11th and 12th ionisation energies. In practice only large jumps such as this are visible on such a graph. The relative values of successive ionisation energies are therefore a direct indicator of the electronic configuration of the atom in question. The trends can be summarised as follows: 1. The successive ionisation energies of an atom always increase. The more electrons that are removed, the fewer the number electrons that remain. There is therefore less repulsion between the electrons in the resulting ion. The electrons are therefore more stable and harder to remove. 2. By far the largest jumps between successive ionisation energies come when the electron is removed from an inner shell. This causes a large drop in shielding, a large increase in effective nuclear charge and a large increase in ionisation energy By applying the above principles in reverse, it is also possible to predict the electronic structure of a species by analysis of the successive ionisation energy data: Eg Si:
ionisation energy
Large jumps occur between 4th and 5th and between 12th and 13th. Therefore there are three shells: The first contains 2 electrons, the second 8 and the third 4.
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23
Topic 1.2
AMOUNT OF SUBSTANCE
The mole Reacting masses and atom economy Solutions and titrations The ideal gas equation Empirical and molecular formulae Ionic equations
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24
THE MOLE
Since atoms are so small, any sensible laboratory quantity of substance must contain a huge number of atoms: 1 litre of water contains 3.3 x 1025 molecules. 1 gram of magnesium contains 2.5 x 1022 atoms. 100 cm3 of oxygen contains 2.5 x 1021molecules. Such numbers are not convenient to work with, so it is necessary to find a unit of "amount" which corresponds better to the sort of quantities of substance normally being measured. The unit chosen for this purpose is the mole. The number is chosen so that 1 mole of a substance corresponds to its relative atomic/molecular/formula mass measured in grams. A mole is thus defined as follows: A mole of a substance is the amount of that substance that contains the same number of elementary particles as there are carbon atoms in 12.00000 grams of carbon-12. One mole of carbon-12 has a mass of 12.0g. One mole of hydrogen atoms has a mass of 1.0g. One mole of hydrogen molecules has a mass of 2.0g. One mole of sodium chloride has a mass of 58.5g. The number of particles in one mole of a substance is 6.02 x 1023. This is known as Avogadro's number, L. Thus when we need to know the number of particles of a substance, we usually count the number of moles. It is much easier than counting the number of particles. The number of particles can be calculated by multiplying the number of moles by Avogadros number. The number of moles can be calculated by dividing the number of particles by Avogadros number.
particles moles L
The mass of one mole of a substance is known as its molar mass, and has units of gmol-1. It must be distinguished from relative atomic/molecular/formula mass, which is a ratio and hence has no units, although both have the same numerical value.
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25 The symbol for molar mass of compounds or molecular elements is m r. The symbol for molar mass of atoms is ar. Mass (m), molar mass (mr or ar) and number of moles (n) are thus related by the following equation:
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26 REACTING MASSES It is possible to use the relationship moles = mass/m r to deduce the masses of reactants and products that will react with each other. When performing calculations involving reacting masses, there are two main points which must be taken into account: The total combined mass of the reactants must be the same as the total combined mass of the products. This is known as the law of conservation of mass. The ratio in which species react corresponds to the number of moles, and not their mass. Masses must therefore all be converted into moles, then compared to each other, then converted back. i) Reactions which go to completion Eg. What mass of aluminium will be needed to react with 10 g of CuO, and what mass of Al2O3 will be produced? 3CuO(s) + 2Al(s) Al2O3(s) + 3Cu(s) 10 g = 10/79.5 = 0.126 moles of CuO 3:2 ratio with Al so 2/3 x 0.126 = 0.0839 moles of Al, so mass of Al = 0.0839 x 27 = 2.3 g 3:1 ratio with Al2O3 so 1/3 x 0.126 = 0.0419 moles of Al2O3, so mass of Al2O3 = 0.0419 x 102 = 4.3 g
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27 ii) Reactions which do not go to completion Many inorganic reactions go to completion. Reactions which go to completion are said to be quantitative. It is because the reactions go to completion that the substances can be analysed in this way. Some reactions, however, particularly organic reactions, do not go to completion. It is possible to calculate the percentage yield of product by using the following equation: % yield = amount of product formed maximum amount of product possible x 100
Eg 2.0 g of ethanol (C2H5OH) is oxidised to ethanoic acid (CH3COOH). 1.9 g of ethanoic acid is produced. What is the percentage yield? (assume 1:1 ratio) Moles of ethanol = 2/46 = 0.0435 Max moles of ethanoic acid = 0.0435 so max mass of ethanoic acid = 0.0435 x 60 = 2.61 g percentage yield = 1.9/2.61 x 100 = 73%
Eg When propanone (CH3COCH3) is reduced to propan-2-ol (CH3CH2CH2OH), a 76% yield is obtained. How much propan-2-ol can be obtained from1.4 g of propanone? (assume 1:1 ratio) Moles of propanone = 1.4/58 = 0.0241 moles So max moles of propan-2-ol produced = 0.0241 moles So actual amount produced = 0.0241 x 76/100 = 0.0183 moles So mass of propan-2-ol = 0.0183 x 60 = 1.1 g
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28 ATOM ECONOMY When we carry out a chemical reaction in order to make a product, we often make other products, called by-products, as well. Eg In the production of NaOH from NaCl the following reaction takes place: 2NaCl + 2H2O 2NaOH + H2 + Cl2 The atom economy of a reaction is the percentage of the total mass of reactants that can, in theory, be converted into the desired product. It can be calculated as follows: % atom economy = mass of desired product x 100 total mass of products
Assuming we start with 2 moles of NaCl and 2 moles of H2O, we will make 2 moles of NaOH, and 1 mole of H2 and Cl2. So % atom economy = (2 x 40) x 100 = (2 x 40) + (1 x 2) + (1 x 71) 52.3 %
The remaining 47.7% of the mass is converted into less useful products and is hence wasted. So the higher the atom economy, the less waste and the more efficient the product process (assuming the reaction does actually go to completion). All reactions which have only one product have an atom economy of 100% Atom economy is an important consideration when considering how to make a particular useful product.
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29 SOLUTIONS A solution is a homogeneous mixture of two or more substances in which the proportions of the substances are identical throughout the mixture. The major component of a solution is called the solvent and the minor components are called the solutes. In most cases water is the solvent. The amount of solute present in a fixed quantity of solvent or solution is called the concentration of the solution. It is usually measured in grams of solute per dm 3 of solution or in moles of solute per dm3 of solution. In the latter case (moldm-3) it is also known as the molarity of the solution. The number of moles of solute, molarity of the solution and volume of solution can thus be related by the equation:
moles
volume(dm3)
molarity
The volume of solution in this case must always be measured in dm 3 (or litres). If the volumes are given in cm3 then V/1000 must be used instead. If concentration is given in gdm-3, it must be converted to molarity before it can be used in the above equation. This can be done easily by dividing by the molar mass of the solute.
C(grams) mr molarity
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30 The volume of one solution required to react with a known volume of another can be deduced from the above relationships and knowledge of the relevant chemical equation. Remember it is moles which react in the ratio shown, so all quantities must be converted to moles before the comparison can be made. The quantitative investigation of chemical reactions by comparing reacting volumes is known as volumetric analysis. The procedure by which reacting volumes are determined is known as a titration. In titrations, a solution whose concentration is unknown is titrated against a solution whose concentration is known. The solution of known concentration is always placed in the burette, and the solution of unknown concentration is always placed in the conical flask. Eg 28.3 cm3 of a 0.10 moldm-3 solution of NaOH was required to react with 25 cm 3 of a solution of H2SO4. What was the concentration of the H2SO4 solution? Equation: H2SO4 + 2NaOH Na2SO4 + 2H2O Moles of NaOH = 28.3/1000 x 0.1 = 2.8 x 10-3 2:1 ratio so moles of H2SO4 = 2.8 x 10-3/2 = 1.4 x 10-3 so concentration of H2SO4 = 1.4 x 10-3/25 x 1000 = 0.056 moldm-3. Eg Calculate the volume of 0.50 moldm-3 nitric acid required to react completely with 5 g of lead (II) carbonate. Equation: PbCO3 + 2HNO3 Pb(NO3)2 + CO2 + H2O Moles of PbCO3 = 5/267 = 0.0187 1:2 ratio so moles of HNO3 = 0.0187 x 2 = 0.0375 Volume of HNO3 = 0.0375/0.5 x 1000 = 74.9 cm3.
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31 GASES
The volume occupied by a gas depends on a number of factors: i) ii) iii) the temperature: the hotter the gas, the faster the particles are moving and the more space they will occupy the pressure: the higher the pressure, the more compressed the gas will be and the less space it will occupy the amount of gas: the more gas particles there are, the more space they will occupy
The volume occupied by a gas does not depend on what gas it is, however: one mole of any gas, at the same temperature and pressure, will have the same volume as one mole of any other gas. The pressure, temperature, volume and amount of gas can be related by a simple equation known as the ideal gas equation:
PV = nRT P is the pressure measured in pascals (Pa) or Nm-2. One atmosphere, which is normal atmospheric pressure, is 101325 Pa. V is the volume in m3. Remember; 1 m3 = 1000 dm3 = 106 cm3. T is the absolute temperature, measured in Kelvin (K). Remember; 0 oC = 273 K. R is the molar gas constant and has a value of 8.31 Jmol-1K-1.
This equation can be rearranged to find the density of gases and the RMM of gases, using the relationship m = n x mr. PV = mRT/mr, so the mass of one mole is given by m r = mRT/PV, where m is the mass in kg. The answer m will also be in kg so it must be converted into grams. The density of a gas, or mass/volume, is given by (m/V) = mrP/RT.
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32 SUMMARY USING MOLES Using the four relationships described, it is possible to calculate the amount of any substance in a chemical reaction provided that the chemical equation is known and the amount of one of the reacting species is also known. The procedure is summarised in the table below:
for gases: n = PV RT for solutions: n = CV
use the ratios in the equation to find the number of moles of other species
n = mass RMM
n = particles L
These relationships are frequently used in practical chemistry. Typical calculations in AS Practical Chemistry involve: i) ii) iii) Determining the concentration of a solution Determining the relative molecular mass of a solid Determining the percentage purity of a solid
The percentage purity of a substance can be calculated as follows: Percentage purity = mass substance would have if it was pure mass of impure substance x 100
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33 EMPIRICAL AND MOLECULAR FORMULAE The empirical formula of a compound is the formula which shows the simplest whole-number ratio in which the atoms in that compound exist. It can be calculated if the composition by mass of the compound is known. The molecular formula of a substance is the formula which shows the number of each type of atom in the one molecule of that substance. It applies only to molecular substances, and can be deduced if the empirical formula and molar mass of the compound are known. The molecular formula is always a simple whole number multiple of the empirical formula. Eg a substance contains 85.8% carbon and 14.2% hydrogen, what is its empirical formula? If its relative molecular mass is 56, what is its molecular formula? Mole ratio = 85.8 12 = 7.15 7.15 1 : 14.2 1 14.2 7.15 2 so empirical formula = CH2
: : :
RMM = 56 = (CH2) so 14n = 56 and n = 56/14 = 4 Molecular formula = C4H8 It is also possible to calculate the percentage composition by mass of a substance, if its empirical or molecular formula is known. Eg What is the percentage composition by mass of ethanoic acid, C2H4O2? RMM = 60 % C = (12 x 2)/60 x 100 = 40.0% % H = (1 x 4)/60 x 100 = 6.67% %O = (16 x 2)/60 x 100 = 53.3%
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34 FORMULAE OF IONIC COMPOUNDS An ion is a species in which the number of electrons is not equal to the number of protons. An ion thus has an overall charge, characteristic of the difference in the number of protons and electrons. Ions with a positive charge are known as cations and ions with a negative charge are known as anions. Compounds made up of ions are known as salts. They are all electrically neutral, so must all contain at least one anion and at least one cation. Salts do not have molecular formulae, as they do not form molecules. They are written as unit formulae. The unit formula of an ionic compound is the formula which shows the simplest whole number ratio in which the ions in the compound exist. This depends on the charges of the ions involved. Some important ions and their charges are shown below: i) cations Formula Na+ K+ Ag+ H+ NH4+ Cu+ Mg2+ Ca2+ Fe2+ Zn2+ Pb2+ Cu2+ Ni2+ Al3+ Cr3+ Fe3+ Name Sodium Potassium Silver Hydrogen Ammonium Copper(I) Magnesium Calcium Iron(II) Zinc Lead(II) Copper(II) Nickel(II) Aluminium Chromium(III) Iron(III)
Charge +1 +1 +1 +1 +1 +1 +2 +2 +2 +2 +2 +2 +2 +3 +3 +3
Note that some atoms can form more than one stable cation. In such cases it is necessary to specify the charge that is on the cation by writing the charge in brackets after the name of the metal.
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35 ii) anions Charge -1 -2 -2 -1 -1 Formula OHSO42CO32NO3HCO3Name Hydroxide Sulphate Carbonate Nitrate Hydrogencarbonate
CHEMICAL EQUATIONS The purpose of chemistry is essentially to study chemical reactions. In chemical reactions, elements or compounds react with each other to form other elements and/or other compounds. Chemical reactions involve the movement of electrons between different species. The nuclei always remain intact. Every chemical reaction can be represented by a chemical equation. A chemical equation indicates the species involved in the reaction and shows the way in which they react. Every chemical equation must contain three pieces of information: i) the identities of all the reactants and products
The chemical formulae of all the species involved in the reaction should be shown. Any species left unchanged should be left out. Reactants must be written on the left of the arrow and products on the right. Remember that in chemical reactions all the nuclei remain unchanged. Therefore the total number of atoms of each type must be the same on each side of the equation. Atoms themselves cannot be created or destroyed in chemical reactions; only transferred from species to species.
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Atoms, elements and compounds combine with each other in simple whole number ratios, eg 1:1, 1:2, 1:3. The ratio in which the species react and in which products are formed are shown in the reaction coefficients. These are the numbers which precede the chemical formula of each species in the equation. If no coefficient is shown it is assumed to be 1. Deducing the reaction coefficients for a reaction is known as balancing the equation. The total number of atoms of each element must be the same on both sides of the equation. When balancing chemical equations, always balance compounds first and elements second. It's easier that way. N.B. Reaction coefficients in no way show the actual amount of a substance which is reacting. They provide information only on the way in which they react. iii) The state symbols
The state symbol shows the physical state of each reacting species and must be included in every chemical equation. There are four state symbols required for Alevel chemistry: (s) - solid (l) - liquid (g) - gas (aq) - aqueous, or dissolved in water
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37
IONIC EQUATIONS
Many reactions that take place in aqueous solution do not involve all of the ions that are written in the equation. Some species remain in aqueous solution before and after the reaction. They therefore play no part in the reaction and are known as spectator ions. In ionic equations, spectator ions are left out. Eg BaCl2(aq) + Na2SO4(aq) BaSO4(s) + 2NaCl(aq) This reaction involves the precipitation of barium sulphate. Notice that the Cl- ions and the Na+ ions remain in the aqueous state before and after the reaction. They are therefore spectator ions. The above reaction can then be rewritten as follows: Ba2+(aq) + SO42-(aq) BaSO4(s) Eg Al2(SO4)3(aq) + 6NaOH(aq) 2Al(OH)3(s) + 3Na2SO4(aq) This reaction involves the precipitation of aluminium hydroxide. The Na+ and SO42- ions are spectator ions and can be left out The ionic equation for the reaction is: Al3+(aq) + 3OH-(aq) Al(OH)3(s) Ionic equations are very useful for simplifying precipitation reactions. They can also simplify acid-base reactions: Eg NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l) The Na+ and Cl- ions are spectator ions, so the ionic equation for the reaction is: H+(aq) + OH-(aq) H2O(l) All reactions between strong acids and strong alkalis have the same ionic equation.
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38
Topic 1.3
BONDING
Types of bond States of matter Structure and physical properties Molecular shapes Intermolecular forces
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39
TYPES OF BOND
Atoms bond to each other in one of four ways:
Ionic Bonding
An ionic bond is an attraction between oppositely charged ions, which are formed by the transfer of electrons from one atom to another. Eg In sodium chloride, each sodium atom transfers an electron to a chlorine atom. The result is a sodium ion and a chloride anion. These two ions attract each other to form a stable compound.
oo Na x o Cl oo o o
+ o x
oo Cl oo o o
Na
+ o x
oo Cl oo o o
Na
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40
Covalent Bonding
A covalent bond is a pair of electrons shared between two atoms. In a normal covalent bond, each atom provides one of the electrons in the bond. A covalent bond is represented by a short straight line between the two atoms. Eg water
x o
x O o x x xx
H O
In a dative covalent bond, one atom provides both electrons to the bond. A dative covalent bond is a pair of electrons shared between two atoms, one of which provides both electrons to the bond. A dative covalent bond is represented by a short arrow from the electron providing both electrons to the electron providing neither. Eg ammonium ion
H+
+ H x o
H H N H H
xx
H
x o
xo
H
Covalent bonding happens because the electrons are more stable when attracted to two nuclei than when attracted to only one. Covalent bonds should not be regarded as shared electron pairs in a fixed position; the electrons are in a state of constant motion and are best regarded more as charge clouds.
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41
Metallic Bonding
A metallic bond is an attraction between cations and a sea of electrons . Metallic bonds are formed when atoms lose electrons and the resulting electrons are attracted to all the resulting cations. Eg Magnesium atoms lose two electrons each, and the resulting electrons are attracted to all the cations.
2+ 2+
Mg
e e e e
Mg
Metallic bonding happens because the electrons are attracted to more than one nucleus and hence more stable. The electrons are said to be delocalized they are not attached to any particular atom but are free to move between the atoms.
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42 IONIC OR COVALENT? - ELECTRONEGATIVITY Electronegativity is the relative ability of an atom to attract electrons in a covalent bond. The electronegativity of an atom depends on its ability to attract electrons and its ability to hold onto electrons. Electronegativity increases across a period as the nuclear charge on the atoms increases but the shielding stays the same, so the electrons are more strongly attracted to the atom. Electronegativity decreases down a group as the number of shells increases, so shielding increases and the electrons are less strongly attracted to the atom. An atom which has a high electronegativity is said to be electronegative; an atom which does not have a high electronegativity is said to be electropositive. Electronegativities are relative; electronegativity has no units and is measured on a scale from 0.7 to 4.0. The electronegativities of some elements in the periodic table are shown below:
H 2.1 Li 1.0 Na 0.9 K 0.8 Be 1.5 Mg 1.2 Ca 1.0 B 2.0 Al 1.5 Ga 1.6 C 2.5 Si 1.8 Ge 1.8 N 3.0 P 2.1 As 2.0 O 3.5 S 2.5 Se 2.4 F 4.0 Cl 3.0 Br 2.8 He Ne Ar Kr
Sc 1.3
Ti 1.5
V 1.6
Cr 1.6
Mn 1.5
Fe 1.8
Co 1.8
Ni 1.8
Cu 1.9
Zn 1.6
Note that the noble gases cannot be ascribed an electronegativity since they do not form bonds. Electronegativity is a very useful concept for predicting whether the bonding between two atoms will be ionic, covalent or metallic. Consider a covalent bond between two atoms A and B.
A x o B
If both atoms have a similar electronegativity, both atoms attract the electrons with similar power and the electrons will remain midway between the two. The bond will thus be covalent - the electrons are shared between the two atoms. Eg H (2.1) and H (2.1)
H x o
H
a covalent bond
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43 If one atom is significantly more electronegative than the other, it attracts the electrons more strongly than the other and the electrons are on average closer to one atom than the other. The electrons are still shared, but one atom has a slight deficit of electrons and thus a slight positive charge and the other a slight surplus of electrons and thus a slight negative charge. Such a bond is said to be polar covalent. Eg H (2.1) and O (3.0)
- x +
O o
symbol respectively. If the difference between the two atoms is large, then the sharing of electrons is so uneven that the more electronegative atom has virtually sole possession of the electrons. The electrons are, in effect, not shared at all but an electron has essentially between transferred from one atom to the other. The more electropositive atom is positively charged and the more electronegative atom is negatively charged. The bonding is thus ionic. Eg Na (0.9) and Cl (3.0) +
Na x Cl o
an ionic bond If both atoms are electropositive, neither has a great ability to attract electrons and the electrons do not remain localised in the bond at all. They are free to move, both atoms gain a positive charge and the bonding is metallic. Eg Mg (1.2) and Mg (1.2)
2+
Mg
x o x o
2+
Mg
a metallic bond
Differences in electronegativity can be used to predict how much ionic or metallic character a covalent bond will have.
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44 Given suitable electronegativity data, it is thus possible to predict whether a bond between two atoms will be ionic, polar covalent, covalent or metallic. If both atoms have electronegativities less than 1.6 - 1.9 then the bond is metallic. If either atom has an electronegativity greater than 1.9 and the difference is less than 0.5 then the bond is covalent. If either atom has an electronegativity greater than 1.9 and the difference is more than 0.5 but less than 2.1 then the bond is polar covalent. If the difference is greater than 2.1 then the bond is ionic.
These rules are not perfect and there are notable exceptions; for example the bond between Si (1.8) and Si (1.8) is covalent but the bond between Cu (1.9) and Cu (1.9) is metallic. The bond between Na (0.9) and H (2.1) is ionic but the bond between Si (1.8) and F (4.0) is polar covalent. However as basic giudelines they are very useful provided that their limitations are appreciated.
All bonds are assumed to be covalent in principle: differences in electronegativity can be used to predict how much ionic or metallic character a covalent bond will have. Electronegativity differences show that bonds between non-identical atoms are all essentially intermediate in character between ionic and covalent. No bond is completely ionic, and only bonds between identical atoms are completely covalent. Bonds between identical atoms cannot be ionic as there is no difference in electronegativity. They will therefore be either covalent or metallic.
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45
STATES OF MATTER
Matter can exist in one of three states; solid, liquid and gas. The state in which a certain substance is most stable at a given temperature depends on the balance between the kinetic energy of the particles, which depends on temperature, and the magnitude of the force of attraction between them.
Solids
In a solid, the particles are tightly packed together in a lattice. A lattice is an ordered and infinitely repeating arrangement of particles. The properties of solids are dominated by the forces in between these particles which cause them to attract each other and preserve this ordered arrangement. A solid thus has a fixed volume and a fixed shape. At all temperatures above absolute zero, the particles have kinetic energy. In a solid, however, this kinetic energy is not enough to cause the particles to fly apart, and nor is it enough to cause significant separation of the particles. The particles are thus restricted to rotational and vibrational motion; no translational motion of the particles with respect to each other is possible. In a solid, the kinetic energy of the particles is not nearly enough to overcome the potential energy caused by their mutual attraction.
SOLIDS
If a solid is heated, the kinetic energy of the particles increases, and they vibrate more. As they vibrate more, the bonds between the particles are weakened, some are broken and spaces appear between the particles. At this point the solid has melted.
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46
Liquids
In a liquid, the particles are by and large packed together in a lattice that extends across the range of 10 - 100 particles. However over a longer range the structure breaks down, and there is enough space between the particles for them to move from one cluster to another. The properties of liquids are still dominated by the forces between the particles, but these particles have enough kinetic energy to move between each other in the spaces that exist. There is thus short-range order but no long-range order. A liquid has a fixed volume but no fixed shape. The kinetic energy of the particles is now significant; it forces the particles apart to the extent that the spaces between them are often wider than the particles themselves. The particles are thus permitted some translational motion with respect to each other within these spaces. All solids will melt if they are heated strongly enough. In a liquid, the kinetic energy of the particles is still not large enough to overcome their mutual attraction, but is nevertheless significant and must be taken into account.
LIQUIDS
Gases
In a gas, all the particles are in rapid and random motion, and thus behave independently of each other. There is no ordered arrangement of any kind, and the spaces between the particles are much larger than the size of the particles themselves. The properties of a gas are dominated by the kinetic energy of the particles; the attraction between them is not significant. A gas has neither a fixed volume nor a fixed shape. In a gas, the kinetic energy of the particles is much greater than the forces of attraction between them. Since the kinetic energy depends only on temperature, it follows that all gases at a similar temperature behave in a similar way. All liquids can be boiled if heated strongly enough.
GASES
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47
IONIC STRUCTURES
Each sodium ion attracts several chloride ions and vice versa so the ionic bonding is not just between one sodium and one chloride ion. There is a 3-D lattice.
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48
+ Na + Na + Na
Cl Cl -
Cl -
Cl
+ Na
+ Na
+ Na
Cl
Cl -
Cl
+ Na
1. Melting and boiling point The attraction between opposite ions is very strong. A lot of kinetic energy is thus required to overcome them and the melting point and boiling point of ionic compounds is very high. In the liquid state, the ions still retain their charge and the attraction between the ions is still strong. Much more energy is required to separate the ions completely and the difference between the melting and boiling point is thus large. Compound Melting point/oC Boiling point/oC NaCl 801 1459 MgO 2852 3600
The higher the charge on the ions, and the smaller they are, the stronger the attraction between them will be and the higher the melting and boiling points. In MgO, the ions have a 2+ and 2- charge and thus the attraction between them is stronger than in NaCl, so the melting and boiling points are higher.
2. Electrical Conductivity Since ionic solids contain ions, they are attracted by electric fields and will, if possible, move towards the electrodes and thus conduct electricity. In the solid state, however, the ions are not free to move since they are tightly held in place by each other. Thus ionic compounds do not conduct electricity in the solid state. Ionic solids are thus good insulators. In the liquid state, the ions are free to move and so can move towards their respective electrodes. Thus ionic compounds can conduct electricity in the liquid state. 3. Mechanical properties Since ions are held strongly in place by the other ions, they cannot move or slip over each other easily and are hence hard and brittle.
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49
METALLIC STRUCTURES
Bonding in metals
Metallic bonding is the attraction between cations and a sea of delocalised electrons. The cations are arranged to form a lattice, with the electrons free to move between them. The structure of the lattice varies from metal to metal, and they do not need to be known in detail. It is possible to draw a simplified form of the lattice: Example - magnesium
2+
e
Mg
Mg
Mg
2+
Mg
2+
ee
Mg
2+
ee
Mg
e
2+
2+
ee
Mg
2+
ee
2+
Mg
Mg
2+
Properties of metals a) Electrical conductivity: since the electrons in a metal are delocalised, they are free to move throughout the crystal in a certain direction when a potential difference is applied and metals can thus conduct electricity in the solid state. The delocalised electron system is still present in the liquid state, so metals can also conduct electricity well in the liquid state. b) Melting and boiling point: although not generally as strong as in ionic compounds, the bonding in metals is relatively strong, and as a result the melting and boiling points of metals are relatively high. Metal Melting point/ oC Boiling point/ oC Na 98 883 K 64 760 Be 127 8 297 0 Mg 649 110 7
Smaller ions, and those with a high charge, attract the electrons more strongly and so have higher melting points than larger ions with a low charge. Na has smaller cations than K so has a higher melting and boiling point. Mg cations have a higher charge than Na so has a higher melting and boiling point. c) Other physical properties: Since the bonding in metals is non-directional, it does not really matter how the cations are oriented relative to each other. The metal cations can be moved around and there will still be delocalized electrons available
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50 to hold the cations together. The metal cations can thus slip over each other fairly easily. As a result, metals tend to be soft, malleable and ductile.
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51
COVALENT STRUCTURES
A covalent bond is a shared pair of electrons between two atoms. When a covalent bond is formed, two atomic orbitals overlap and a molecular orbital is formed. Like atomic orbitals, a molecular orbital can only contain two electrons. Overlap of atomic orbitals is thus only possible if both orbitals contain only one electron (normal covalent bond), or if one is full and the other empty (dative covalent bond). Covalent bonding happens because the electrons are more stable when attracted to two nuclei than when attracted to only one:
Normal covalent bonds An overlap between two orbitals, each containing one electron, is a normal covalent bond. The number of normal covalent bonds which an atom can form depends on its number of unpaired electrons. Some atoms, like carbon, promote electrons from s to p orbitals to create unpaired electrons. 1s 2s 2p F has 1 unpaired electron in a 2p orbital forms one covalent bond
xx x x
F
xx
x o
Eg hydrogen fluoride 1s 2s
2p
Carbon rearranges slightly to make more unpaired electrons 1s 2s 2p C has 4 unpaired electrons forms four covalent bonds
H
ox x o x o
C
xo
Eg methane
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52 Dative covalent bonds Any atom which has filled valence shell orbitals can provide both electrons for a dative covalent bond. This includes any element in groups V, VI, VII or 0 but is most common in N, O and Cl. 1s 2s 2p N has three unpaired electrons and one electron pair
Any atom which has empty valence shell orbitals can accept a pair of electrons for a covalent bond. This includes any element in groups I, II and III but is most common in Be, B and Al. 1s 2s 2p
B promotes an electron from 2s to 2p to form 3 unpaired electrons: 1s 2s 2p B has 3 unpaired electrons and an empty orbital
H
ox x o x x
H
ox
B
xo
x o xo
Eg BH3NH3 Sigma and pi bonds Atomic orbitals can overlap in one of two ways: If they overlap directly along the internuclear axis, as is most common, a -bond is formed.
or
A -bond is a bond resulting from direct overlap of two orbitals along the internuclear axis. All single bonds between two atoms are -bonds. It is only possible to form one -bond between two atoms, since another would force too many electrons into a small space and generate repulsion. If double bonds are formed, therefore, the orbitals must overlap in a different way.
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53
If two orbitals overlap above and below (or behind and in front of) the internuclear axis, then a -bond is formed.
or
A -bond is a bond resulting from overlap of atomic orbitals above and below the internuclear axis. All double bonds consist of a -bond and a -bond. All triple bonds consist of a -bond and two -bonds. If the first -bond results from overlap above and below the internuclear axis, the second results from overlap behind and in front of the internuclear axis. Note that -bonds can only be formed by overlap of p-orbitals, since s-orbitals do not have the correct geometry. -bonds can also be represented by orbital diagrams. Eg ethene:
Strength of covalent bonds Covalent bonds are in general strong. The smaller the atoms, the closer the electrons are to the two nuclei and the stronger the bond. Bond C-F C-Cl C-Br C-I Bond dissociation energy/ kJmol-1 467 346 290 228
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54
I I I I... I I ... I
There are attractive forces between these molecules, known as intermolecular forces, but they are weak. In the gaseous state, the intermolecular forces are broken but the bonds within the molecule remain intact - they are not broken. The gas phase consists of molecules, not atoms. Molecular substances have certain characteristic properties: Melting and boiling point: these are generally low, since intermolecular forces are weak. Intermolecular forces also decrease rapidly with increasing distance, so there is often little difference in the melting and boiling points. Substance Melting point /oC Boiling point /oC CH4 -184 -166 H2O 0 100 H2 -259 -253 He -272 -268
Electrical conductivity: There are no ions and no delocalised electrons, so there is little electrical conductivity in either solid or liquid state. Other physical properties: The intermolecular forces are weak and generally nondirectional, so most molecular covalent substances are soft, crumbly and not very strong.
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55 Giant covalent In some cases, it is not possible to satisfy the bonding capacity of a substance in the form of a molecule; the bonds between atoms continue indefinitely, and a large lattice is formed. There are no discrete molecules and covalent bonding exists between all adjacent atoms. Such substances are called giant covalent substances, and the most important examples are C, B, Si and SiO2. Example diamond (diamond is an allotrope of carbon) Don't forget that this is just a tiny part of a giant structure extending on all 3 dimensions.
In giant covalent compounds, covalent bonds must be broken before a substance can melt or boil. Giant covalent compounds have certain characteristic properties: Melting and boiling point: these are generally very high, since strong covalent bonds must be broken before any atoms can be separated. The melting and boiling points depend on the number of bonds formed by each atom and the bond strength. The difference between melting and boiling points is not usually very large, since covalent bonds are very directional and once broken, are broken completely. Substance Melting point /oC Boiling point /oC C 3550 4827 Si 1410 2355 B 2300 2550 SiO2 1510 2230
Electrical conductivity: there are no ions or delocalised electrons, so there is little electrical conductivity in either solid or liquid state. Other physical properties: since the covalent bonds are strong and directional, giant covalent substances are hard, strong and brittle. Diamond is in fact the hardest substance known to man. For this reason it is used in drills, glass-cutting and styluses for turntables.
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56 Giant covalent layered Some substances contain an infinite lattice of covalently bonded atoms in two dimensions only to form layers. The different layers are held together by intermolecular forces, and there are often delocalized electrons in between the layers. Examples of these structures are graphite and black phosphorus.
Example - graphite
or In graphite, each carbon atom is bonded to three others. The spare electron is delocalized and occupies the space in between the layers. All atoms in the same layer are held together by strong covalent bonds, and the different layers are held together by intermolecular forces. A number of characteristic properties of graphite result from this structure: Electrical conductivity: due to the delocalised electrons in each plane, graphite is a very good conductor of electricity in the x and y directions, even in the solid state (unusually for a non-metal). However, since the delocalisation is only in two dimensions, there is little electrical conductivity in the z direction (i.e. perpendicular to the planes). Density: graphite has a much lower density than diamond (2.25 gcm -3) due to the relatively large distances in between the planes. Hardness: graphite is much softer than diamond since the different planes can slip over each other fairly easily. This results in the widespread use of graphite in pencils and as an industrial lubricant.
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57
Nature of bonding
Attraction between oppositely charged ions. Infinite lattice of oppositely charged ions in three dimensions
Physical properties
High mpt, bpt Good conductors in liquid state Poor conductors in solid state Hard, strong, brittle
METALLIC Eg Mg
Attraction between cations and delocalised electrons. Infinite lattice of cations in three dimensions, with delocalized electrons in the spaces Infinite lattice of atoms linked by covalent bonds in three dimensions. Covalent bonds are pairs of electrons shared between two atoms Discrete molecules. Atoms in molecule linked by covalent bonds. Weak intermolecular between molecules. forces
High mpt, bpt Good conductors in solid state Good conductors in liquid state Strong, malleable
Very high mpt, bpt Poor conductors in solid state Poor conductors in liquid state Hard, strong, brittle
MOLECULAR Eg I2
Low mpt, bpt Poor conductors in solid state Poor conductors in liquid state Soft, weak, powdery
GIANT COVALENT Infinite lattice of atoms linked by High mpt, bpt covalent bonds in two dimensions Good conductors parallel to planes LAYERED Eg graphite
to form planes. Planes held together by intermolecular forces. Delocalised electrons in between layers Poor conductors perpendicular to planes Soft
Dont forget to learn the structures of Sodium chloride Iodine Diamond Graphite
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MOLECULAR SHAPES
When an atom forms a covalent bond with another atom, the electrons in the different bonds and the non-bonding electrons in the outer shell all behave as negatively charged clouds and repel each other. In order to minimise this repulsion, all the outer shell electrons spread out as far apart in space as possible. Molecular shapes and the angles between bonds can be predicted by the VSEPR theory VSEPR = valence shell electron pair repulsion VSEPR theory consists of two basic rules: i) ii) All -bonded electron pairs and all lone pairs arrange themselves as far apart in space as is possible. -bonded electron pairs are excluded. Lone pairs repel more strongly than bonding pairs.
These two rules can be used to predict the shape of any covalent molecule or ion, and the angles between the bonds. a) 2 electron pairs If there are two electron pairs on the central atom, the angle between the bonds is 180o.
Molecules which adopt this shape are said to be LINEAR. E.g. BeCl2, CO2 b) three electron pairs If there are three electron pairs on the central atom, the angle between the bonds is 120o.
Molecules which adopt this shape are said to be TRIGONAL PLANAR. E.g. BF3, AlCl3, CO32-, NO3-
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59 If one of these electron pairs is a lone pair, the bond angle is slightly less than 120 o due to the stronger repulsion from lone pairs, forcing them closer together.
bond angle = 118o Molecules which adopt this shape are said to be BENT. E.g. SO2, NO2c) Four electron pairs If there are four bonded pairs on the central atom, the angle between the bonds is approx 109o.
Molecules which adopt this shape are said to be TETRAHEDRAL. E.g. CH4, SiCl4, NH4+, SO42If one of the electron pairs is a lone pair, the bond angle is slightly less than 109 o, due to the extra lone pair repulsion which pushes the bonds closer together (approx 107o).
Molecules which adopt this shape are said to be TRIGONAL PYRAMIDAL. E.g. NH3, PCl3
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60 If two of the electron pairs are lone pairs, the bond angle is also slightly less than 109o, due to the extra lone pair repulsion (approx 104o).
Molecules which adopt this shape are said to be BENT. E.g. H2O, OF2 d) Five electron pairs If there are five bonded pairs on the central atom, the three bonds are in a plane at 120o to each other, the other 2 are at 90 o to the plane.
d) Six electron pairs If there are six electron pairs on the central atom, the angle between the bonds is 90o.
Molecules which adopt this shape are said to be OCTAHEDRAL. E.g. SF6
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61 If there are 4 bonding pairs and 2 lone pairs, the bonded pairs are at 90 o in the plane and the lone pairs at 180o. The angles are still exactly 90o because the lone pairs are opposite each other so their repulsion cancels out.
Molecules which adopt this shape are said to be SQUARE PLANAR. E.g. XeF4, ClF4SUMMARY OF MOLECULAR SHAPES Valence shell Bonding Lone shape electron pairs pairs pairs 2 2 0 LINEAR A B B 3 3 0 TRIGONAL PLANAR B B
A B
BENT B
115 - 118
A B
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62 4 4 0 TETRAHEDRAL B
A B B
109.5
TRIGONAL PYRAMIDAL
107
A B
4 2 2
B BENT
A B B
104.5
TRIGONAL BIPYRAMIDAL B
90 and 120
B B A B
6 6 0
90
90
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63
INTERMOLECULAR FORCES
There are no covalent bonds between molecules in molecular covalent compounds. There are, however, forces of attraction between these molecules, and it is these which must be overcome when the substance is melted and boiled. These forces are known as intermolecular forces. There are three main types of intermolecular force: 1. Van der Waal's forces Consider a molecule of oxygen, O2.
e eOe e e e e e e Oe e
The electrons in this molecule are not static; they are in a state of constant motion. It is therefore likely that at any given time the distribution of electrons will not be exactly symmetrical - there is likely to be a slight surplus of electrons on one of the atoms.
+ eOe
e e e Oe e ee e e
This is known as a temporary dipole. It lasts for a very short time as the electrons are constantly moving. Temporary dipoles are constantly appearing and disappearing. Consider now an adjacent molecule. The electrons on this molecule are repelled by the negative part of the dipole and attracted to the positive part, and move accordingly.
+ eOe
e e e Oe e ee e e
+ eOe
e e e Oe e ee e e
This is known as an induced dipole. There is a resulting attraction between the two molecules, and this known as a Van der Waal's force. Van der Waal's forces are present between all molecules, although they can be very weak. They are the reason all compounds can be liquefied and solidified. Van der Waal's forces tend to have strengths between 1 kJmol-1 and 50 kJmol-1.
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64 The strength of the Van der Waal's forces in between molecules depends on two factors: a) the number of electrons in the molecule
The greater the number of electrons in a molecule, the greater the likelihood of a distortion and thus the greater the frequency and magnitude of the temporary dipoles. Thus the Van der Waal's forces between the molecules are stronger and the melting and boiling points are larger. Eg noble gases: Substance Number of electrons Melting point/oC Boiling point/oC Eg alkanes: Substance Number of electrons Melting point/oC Boiling point/oC CH4 10 -182 -164 C2H6 18 -183 -88 C3H8 26 -190 -42 C4H10 34 -138 0 He 2 -272 -269 Ne 10 -252 -250 Ar 18 -189 -186 Kr 36 -157 -152
a) Surface area of the molecules The larger the surface area of a molecule, the more contact it will have with adjacent molecules. Thus the greater its ability to induce a dipole in an adjacent molecule and the greater the Van der Waal's forces and melting and boiling points. This point can be illustrated by comparing different molecules containing a similar number of electrons: Substance Kr Number of 36 electrons Melting point/oC -157 o Boiling point/ C -152 Cl2 34 -101 -35 CH3CH(CH3)CH3 CH3CH2CH2CH3 34 34 -159 -12 -138 0
Page 64
65 CH3CH(CH3)CH3 methylpropane
H H C H H H C C H H H
CH3CH2CH2CH3 butane
H H C H H C H H C H H
H
C H
C H
Note that butane has a larger surface area than 2-methylpropane, although they have the same molecular formula (C4H10). Straight-chain molecules always have higher boiling points than their isomers with branched chains. 2. Dipole-dipole bonding Temporary dipoles exist in all molecules, but in some molecules there is also a permanent dipole. Most covalent bonds have a degree of ionic character resulting from a difference in electronegativity between the atoms. This results in a polar bond and a dipole.
+ H
Cl
+ C
In many cases, however, the presence of polar bonds (dipoles) does not result in a permanent dipole on the molecule, as there are other polar bonds (dipoles) in the same molecule which have the effect of cancelling each other out. This effect can be seen in a number of linear, trigonal planar and tetrahedral substances: Cl F F
C Cl Cl Cl
B F
C
CO2
BF3 CCl4
In all the above cases, there are dipoles resulting from polar bonds but the vector sum of these dipoles is zero; i.e. the dipoles cancel each other out. The molecule thus has no overall dipole and is said to be non-polar. Non-polar molecules are those in which there are no polar bonds or in which the dipoles resulting from the polar bonds all cancel each other out. The only intermolecular forces that exist between non-polar molecules are temporaryinduced dipole attractions, or Van der Waals forces. In other molecules, however, there are dipoles on the molecule which do not cancel each other out:
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66
+ H
C Cl Cl CHCl3
Cl
O - S
N
+
H H +
O SO2
NH3
In all the above cases, there are dipoles resulting from polar bonds whose vector sum is not zero; i.e. the dipoles do not cancel each other out. The molecule thus has a permanent dipole and is said to be polar. Polar molecules are those in which there are polar bonds and in which the dipoles resulting from the polar bonds do not cancel out. In addition to the Van der Waal's forces caused by temporary dipoles, molecules with permanent dipoles are also attracted to each other by dipole-dipole bonding. This is an attraction between a permanent dipole on one molecule and a permanent dipole on another.
+ H Br + H Br + H Br
Dipole-dipole bonding usually results in the boiling points of the compounds being slightly higher than expected from temporary dipoles alone; it slightly increases the strength of the intermolecular bonding. The effect of dipole-dipole bonding can be seen by comparing the melting and boiling points of different substances which should have Van der Waal's forces of similar strength: Substance Number of electrons Permanent dipole Melting point/oC Boiling point/oC Cl2 34 NO -101 -45 HBr 36 YES -88 -67 CH3CH(CH3)CH3 CH3COCH3 34 32 NO -159 -73 YES -95 -44
3. Hydrogen bonding In most cases as seen above, the presence of permanent dipoles only makes a slight difference to the magnitude of the intermolecular forces. There is one
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67 exceptional case, however, where the permanent dipole makes a huge difference to the strength of the bonding between the molecules. Consider a molecule of hydrogen fluoride, HF. This clearly has a permanent dipole as there is a large difference in electronegativity between H (2.1) and F (4.0). The electrons in this bond are on average much closer to the F than the H:
+
H
The result of this is that the H atom has on almost no electron density around its nucleus at all and is therefore very small. The H atom is therefore able to approach electronegative atoms on adjacent molecules very closely and form a very strong intermolecular dipole-dipole bond.
F This is known as hydrogen bonding. It is only possible if the hydrogen atom is bonded to a very electronegative element; i.e. N, O or F. It is not fundamentally different from dipole-dipole bonding; it is just a stronger form of it.
H H H
A hydrogen bond can be defined as an attraction between an electropositive hydrogen atom (ie covalently bonded to N, O or F) and an electronegative atom on an adjacent molecule. Examples of substances containing hydrogen bonds are HF, H 2O, NH3, alcohols, carboxylic acids, amines, acid amides and urea.
O H H N H H H H H hydrogen bonds N H H H O H H N H H H O
a) Effect on boiling point The effect of hydrogen bonding on melting and boiling points of substances is huge, unlike other dipole-dipole bonds. Many substances containing hydrogen bonds have much higher boiling points than would be predicted from Van der Waal's forces alone.
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68 Substance Structure
CH3
CH3OCH3
O C
CH3
CH3CH2OH
CH3CH2CH2CH O
H H
C
CH3CH2COOH
H H C H C OH H O
H
C
O
C
CH3
CH2
OH
26 NO -95 -44
26 YES -117 79
40 NO -81 56
Another important series of trends are the boiling points of the hydrides of elements in groups V, VI and VII of the periodic table: Group V: NH3, PH3, AsH3, SbH3 Group VI: H2O, H2S, H2Se, H2Te Group VII: HF, HCl, HBr, HI The boiling points of these graphs are shown graphically below:
Graph to show how hydrogen bonding affects boiling point
150
100
H2O
50 HF 0 0 -50 20 NH3 40 60 80 H2Se AsH3 HBr 100 H2Te 120 SbH3 140 HI
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69 In each case the hydride of period 2 shows a boiling point which is abnormally high ( H2O, NH3 and HF). The general increase in boiling point down the groups result from the increase in Van der Waal's forces which results from an increasing number of electrons in the molecules. There are permanent dipoles but they are not very strong. The abnormally high boiling points of H2O, NH3 and HF are a result of hydrogen bonding between the molecules. Thus results in very strong intermolecular forces between the molecules despite the fact that the Van der Waal's forces are weaker than in the other hydrides. b) other effects of hydrogen bonding The effects of hydrogen bonding on the physical properties of a substance are not restricted to elevated melting and boiling points; it can influence the properties of substances in other ways: The low density of ice. This is due to hydrogen bonding. In ice, the water molecules arrange themselves in such a way as to maximise the amount of hydrogen bonding between the molecules. This results in a very open hexagonal structure with large spaces within the crystal. This accounts for its low density. When the ice melts, the structure collapses into the open spaces and the resulting liquid, despite being less ordered, occupies less space and is thus more dense. Thus ice floats on water. The helical nature of DNA. This is also due to hydrogen bonding. Molecules of DNA contain N-H bonds and so hydrogen bonding is possible. The long chains also contain C=O bonds and the H atoms can form a hydrogen bond with this electronegative O atom. This results in the molecule spiralling, as the C=O bonds and the N-H bonds approach each other. This is an example of an intramolecular hydrogen bond, where the attraction is between a hydrogen atom and an electronegative atom on the same molecule. This must be distinguished from intermolecular hydrogen bonding, in which the attraction is between a hydrogen atom and an electronegative atom on an adjacent molecule.
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70
Topic 1.4
PERIODICITY
The Periodic Table Trends in Period 3 Trends in Group II
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71
II
H Be Mg Ca Sr Ba Ra
III IV
B Al Ga In Tl C Si Ge Sn Pb
V VI VII
N P As Sb Bi O S Se Te Po F Cl Br I At
0
He Ne Ar Kr Xe Rn
Sc Y La Ac
Ce - Lu Th - Lw
Ti Zr Hf
V Nb Ta
Cr Mo W
Mn Tc Re
Fe Ru Os
Co Rh Ir
Ni Pd Pt
Cu Ag Au
Zn Cd Hg
Since the electronic configurations of H and He are unusual, they do not fit comfortably into any group. They are thus allocated a group based on similarities in physical and chemical properties with other members of the group. He is placed in group 0 on this basis, but hydrogen does not behave like any other element and so is placed in a group of its own. The elements Ce - Lu and Th - Lw belong in the periodic table as shown above. However if they are placed there periods 6 and 7 do not fit onto a page of A4, so they are placed below the other elements in most tables.
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72 All elements belong to one of four main blocks: the s-block, the p-block, the d-block and the f-block. The s-block elements are all those with only s electrons in the outer shell. The p-block elements are all those with at least one p-electron in the outer shell. The d-block elements are all those with at least one d-electron and at least one s-electron but no f or p electrons in the outer shell. The f-block elements are all those with at least one f-electron and at least one selectron but no d or p electrons in the outer shell. I
Li Na K Rb Cs Fr
II
H Be Mg Ca Sr Ba Ra
III IV
B Al Ga In Tl C Si Ge Sn Pb
V VI VII 0
He N P As Sb Bi O S Se Te Po F Cl Br I At Ne Ar Kr Xe Rn
Sc Y La Ac
Ce - Lu Th - Lw
Ti Zr Hf
V Nb Ta
Cr Mo W
Mn Tc Re
Fe Ru Os
Co Rh Ir
Ni Pd Pt
Cu Ag Au
Zn Cd Hg
Elements coloured green are in the s-block Elements coloured blue are in the p-block Elements coloured red are in the d-block Elements coloured black are in the f-block The physical and chemical properties of elements in the Periodic Table show clear patterns related to the position of each element in the Periodic Table. Elements in the same group show similar properties, and properties change gradually on crossing a Period. As atomic number increases, the properties of the elements show trends which repeat themselves in each Period of the Periodic Table. These trends are known as Periodic Trends and the study of these trends in known as Periodicity.
Page 72
73 TRENDS IN PERIOD 3
0.2
Na
0.15
nm
Mg
Al Si P S Cl Ar
0.1
0.05
b)
Ionization energies
Ionisation energy generally increases across period 3 but decreases between groups II and III and also between groups V and VI. Ionisation energy increases across period 3 because the nuclear charge increases but the shielding remains the same, making the electrons harder to remove. Ionisation energy decreases from group II to group III because the outer electron in Al is in a 3p orbital, but the outer electron in Mg is in a 3s orbital. The 3p orbital is better shielded from its nucleus making it easier to remove. Ionisation energy decreases from group V to group VI because the outermost 3p electron in S is paired, so there is repulsion in the orbital and the electron is easier to remove. The outermost 3p electron in P is unpaired, so experiences less repulsion and is harder to remove.
Ist ionisation energies of Period 3 elements
1800 1600
Ar
Page 73
74 c) Electronegativity
Electronegativity increases across period 3. As the nuclear charge increases but the shielding remains the same, the electrons are attracted more strongly to the atom, so that atom will have a larger share of the electrons in a covalent bond.
The variation on bond type causes a number of differences in the structures of the Period 3 elements which in turn causes significant differences in physical properties. a) Sodium, Magnesium and Aluminium
Sodium, Magnesium and Aluminium are metals. They consist of an infinite lattice of cations held together by a sea of delocalised electrons. There is a fairly strong attraction between the cations and the delocalised electrons and as a result metals tend to have fairly high melting points and boiling points. The melting points increase with increasing charge and decreasing size and thus increase across a period. Element Mpt/oC Bpt/oC Sodium 98 883 Magnesium 669 1107 Aluminium 680 2467
The delocalised electrons in the metal structure are free to move throughout the metal lattice and can thus behave as charge carriers. When a potential difference is
Page 74
75 applied, the electrons can move towards the positive electrode. Thus metals are good conductors of electricity. Electrical conductivity increases from sodium to aluminium as the number of delocalized electrons per atom increases. Aluminium has three electrons per atom in the sea, magnesium two per atom and sodium only one per atom. b) Silicon
Silicon is a giant covalent macromolecule. Silicon atoms form infinite lattices in which all the atoms are held together by strong covalent bonds. Since the structure cannot be broken up without breaking these strong covalent bonds, it follows that silicon has a very high melting and boiling point. The structure of silicon is tetrahedral, identical to diamond:
Structure
SILICON:
Mpt/oC Bpt/oC
1406 2355
Silicon does not conduct electricity well as it has no free electrons and no free ions. c) Phosphorus, sulphur, chlorine and argon
Phosphorus, sulphur, chlorine and argon form simple molecular structures.There are strong, covalent bonds within the molecule but the different molecules are only held together by weak Van der Waal's forces. Separating these molecules thus requires little energy and the melting and boiling points of these elements are relatively low. The larger the molecule, the greater the magnitude of the temporary and induced dipoles and the higher the melting and boiling points.
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76 Element Structure
P
Phosphorus
P P P P P P P P P P P P P P P
S S S
Sulphur
S S S S S S S S S S S S S S S S S S S S S
Chlorine
Cl Cl Cl Cl Cl Cl Cl Cl Cl Cl
Argon
Ar Ar Ar Ar Ar Ar
44 280 P4 (or P)
-189 -186 Ar
Sulphur has the highest melting point as it exists as S8 molecules. These molecules are quite large, so the number of electrons in the molecule is high and the Van der Waals forces are quite strong. Phosphorus exists as P4 molecules, which have fewer electrons in them and so have weaker Van der Waals forces. So phosphorus has a lower melting point than sulphur. Chlorine exists as molecules, which have even fewer electrons in them so the Van der Waals forces are lower and chlorine has a lower melting point than sulphur and phosphorus. Argon has the lowest melting and boiling point of all, as it exists as single Ar atoms which have even less electrons and so only form very weak Van der Waals forces. These elements do not conduct electricity well as they have no free electrons and no free ions.
Electrical conductivity
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77
Sodium Metallic
Silicon covalent
Sulphur covalent
Chlorine covalent
Argon -
e 3+ 3+ 2+ + + 2+ Mg Na Na Mg Al Al e e e e ee e e e e e e e ee e e e e 3+ e 3+ 2+ 2+ + + Mg Al Mg Al Na Na e ee
P P P P P P P P P
P P P P P P
S S S S S
S S S S S S
S S S
S S S S S S S S S S
Cl Cl
Cl Cl Cl Cl Cl Cl Cl Cl
A Ar Ar Ar Ar
Type: Melting point/oC: Boiling point/oC: First IE/ kJmol-1: Enegativit y Electrical conductiv ity (x10-81 -1 m )
0.21
0.26
0.41
NB You do not need to know the exact figures, just know the trends and be able to explain them.
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78
Topic 1.5
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79
H C H H
H C C H
H C H H H
C H
Branched molecule
Carbon atoms can also be arranged to form rings. These are known as cyclic molecules. The most common number of carbon atoms in a ring is 6.
H H H H H H H C C C H C H C C H H H
Cyclic molecule
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80 b) The functional groups in the molecule A functional group is a specific atom or group of atoms which confer certain physical and chemical properties onto the molecule. Organic molecules are classified by the dominant functional group on the molecule.
Functional groups
These are the some of the most important functional groups found on organic molecules:
Type of compound Alkane Nature of functional group C-C and C-H single bonds only (ie no functional group)
C
Alkene C=C double bond
C
Haloalkane -Chloroalkane -Bromoalkane -Iodoalkane
H H C H
H C H H
H C C H
H C H H H
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81 b) Structural formula, not showing all covalent bonds Enough information is shown to make the structure clear, but most of the actual covalent bonds are omitted. Only important bonds are always shown. Straight chain alkanes are shown as follows:
H H C H H C H H C H H C H H
is represented as CH3CH(CH3)CH3.
H
H H H H C H H H C C C H C H H H
is represented as CH3C(CH3)2CH3.
is represented as CH2=CHCH3.
H
H H C H C H C H
C H
is represented as CH3CH=CHCH3.
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82 c) Molecular formula The molecular formula shows the number of each atom in one molecule of the compound. It does not show unequivocally the structure of the molecule.
H H C H H C H H C H H C H H
is written C4H10
H H C H C H C H
H C H H
is written C4H8
Alkanes have the general molecular formula CnH2n+2. Alkenes have the general molecular formula CnH2n. Haloalkanes have the general molecular formula CnH2n+1X. d) Empirical formula The empirical formula is the simplest whole number ratio of the number of atoms of each element in a substance.
Homologous series
Organic compounds with the same functional group, but a different number of carbon atoms, are said to belong to the same homologous series. Every time a carbon atom is added to the chain, two hydrogen atoms are also added. A homologous series is a series of organic compounds which have the same functional group, but in which the formula of each successive member increases by -CH2-. Eg Homologous series of alkanes: CH4, CH3CH3, CH3CH2CH3, CH3CH2CH2CH3, CH3(CH2)3CH3, CH3(CH2)4CH3 etc As a homologous series is ascended, the size of the molecule increases. Therefore the Van der Waals forces between the molecules become stronger and the boiling point increases.
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83 NOMENCLATURE OF ORGANIC COMPOUNDS Most organic compounds can be named systematically by the IUPAC method. In order to describe completely an organic molecule, three features must be described: the longest straight carbon chain on the molecule. the length and position of any branches on the molecule. the nature and position of any functional groups on the molecule.
The longest straight chain on the molecule The longest straight chain on the molecule is indicated by one of the following prefixes:
Number of carbon atoms in the chain 1 2 3 4 5 6 MethEthPropButPentHexPrefix
Alkanes
Alkanes are named using the ending -ane:
H H C H
H H C H
H H C H
Methane
H C H
H C H H C H H
Ethane
Propane
H H C H
H C H
H C H
H C H H
Butane
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84
Alkenes
Alkenes are named using the ending -ene. In molecules with a straight chain of 4 or more carbon atoms, the position of the C=C double bond must be specified. The carbon atoms on the straight chain must be numbered, starting with the end closest to the double bond. The lowest-numbered carbon atom participating in the double bond is indicated just before the -ene:
H C H
H H C H H
H H C H H
H H C H C H C H
H C H
H C H
Ethene
Propene
H C
H C H
H C H H
But-1-ene
But-2-ene
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85
Haloalkanes
Haloalkanes are named using the prefix chloro-, bromo- or iodo-, with the ending ane. In molecules with a straight chain of three or more carbon atoms, the position of the halogen atom must also be specified. The carbon atoms on the straight chain must be numbered, starting with the end closest to the halogen atom. The number of the carbon atom attached to the halogen is indicated before the prefix:
H H C H
H H C H H C Br
Chloroethane
C H
H C H H
Cl
2-bromopropane
H I C H
H H C H
H C H
H C H
H C H
Cl C H
H C H
H C H
H C H
H C H H
1-iodopentane
3-chloropentane
The position of all halogens in dihaloalkanes except those with one carbon atom must be specified. If there is more than one of the same type of halogen atom on the molecule, the di (two), tri (three) or tetra (four) prefixes must also be used.
H Cl C Cl
H Cl C H
H Br C H
1,1-dichloroethane
C H
H C Cl
H C Cl H C H H
1,2-dichloroethane
1-bromo,2chloropropane
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86
Prefix Methyl
Ethyl
C H
The position of the branch must be specified according to the number of the carbon on the straight chain to which it is attached. The carbons are always numbered from the carbon at the end of the chain closest to the functional group. If there is no functional group, the carbons are numbered from the carbon at the end of the chain closest to the branch. Eg
H H H C C C H H H C H H H
H H C H
H C H H
H C C H
H C H H H
H
C H H
2-methylbutane
2,2-dimethylpropane
H H H H H C H H C H C C H H H C H H H C C
H H H H C H H H
H C C C C C H H H H H C H H H C H H
H
H C H
C H H
2-methyl,3-ethylpentane
3,3-diethylpentane
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87
CH3
CH CH3
CH2
Cl
2-methyl,2-chloropropane
2-methyl,1-chloropropane
Many organic compounds which appear to be different are in fact the same. They appear to be different because different notations are used, or because some of the bonds are simply rotated. Eg butane can be represented in a number of ways: Such as
H
H
H H C H H C H H C H H C H H
H C H H
H C C H H H
H H
C C H H
H H C C H H H
C H
H H C Cl
H C H
H C H H
Cl
H C H
H C H
H C H H
Cl
C H H
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88
ISOMERISM
Isomers are molecules which have the same molecular formula but different structures. There are a number of different types of isomerism in organic compounds, but the only type required for AS Chemistry is structural isomerism. Structural isomers are molecules which have the same molecular formula but a different arrangement of covalent bonds. The different arrangement of covalent bonds can result from: i) The functional group being in different positions (positional isomerism) ii) A different arrangement of the carbon skeleton (chain isomerism) iii) A different functional group (functional isomerism) i) Positional isomerism
Positional isomers are molecules with the same molecular formula but which have the functional group on different positions in the molecule. Alkanes do not show functional isomerism as they have no functional group. Alkenes with four or more carbon atoms show positional isomerism: Eg but-1-ene and but-2-ene
H H C H H H C C C H H H
H H C H C H C H
H C H H
Haloalkanes with three or more carbon atoms show positional isomerism Eg 1-chloromethylpropane and 2-chloromethylpropane
Cl
CH3 CH CH3 CH2 Cl
CH3
C CH3
CH3
ii)
Chain isomerism
Chain isomers are molecules with the same molecular formula but a different arrangement of carbon atoms. The arrangement of carbon atoms in an organic molecule is known as the carbon skeleton.
Page 88
89 Carbon skeletons containing up to three carbon atoms can only be arranged in one way i.e. a straight chain with no branching:
Carbon skeletons containing four carbon atoms can be arranged in two ways:
C
C C C C
C C
Carbon skeletons containing five carbon atoms can be arranged in three ways:
C C C C C C C Carbon skeletons containing six carbon atoms can be arranged in five ways:
C C C C C C
C C
C
C
C
C
C C
C C
All molecules containing four or more carbon atoms can thus show chain isomerism: Eg. Butane and methylpropane
H
H H C H H C H H C H H C H H
H C C H
H C H H H
C H H
Page 89
C H C H C
H H C H
H C H H
H H
C
H H
C
H
C
H
C
H
C
H
C
Cl
Cl
H H
C H H
iii)
Functional isomerism
Functional isomers are molecules with the same molecular formula but different functional groups. eg Alkanes which have a ring rather than a straight chain arrangement are known as cycloalkanes. They have the general formula CnH2n, which is the same as alkenes. Cycloalkanes and alkenes can thus show functional isomerism. Eg cyclohexane and hex-1-ene
H H H H H H H C C C H C H C C H H
H H C H H H H H C H C H C H C H C H
H H C H C
H C H
H C H H
Page 90
Isomers tend to differ slightly in their melting and boiling points. Molecules with no branching tend to have higher boiling points than isomers with more branching. This is because they have a higher surface area, so they pack together better and so the van der Waals forces are stronger. Eg isomers of pentane, C5H12: Isomer Pentane Structure
H H C H H C H H C H H C H H C H H
Boiling point/oC
36
Methylbutane
H
H C H
H C H H
H C C H
H C H H H
28
2,2-dimethylpropane
H H H
C
H H
C C
H
C
10
H H
C H
H H
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92
Topic 1.6
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93
CRUDE OIL
1. Introduction The vast majority of carbon-containing compounds in widespread use have been made from crude oil. Crude oil is also known as petroleum. Crude oil is a mixture of hydrocarbons. A hydrocarbon is a substance containing carbon and hydrogen only. Most of the hydrocarbons in crude oil are alkanes. Alkanes are hydrocarbons containing only single bonds between the carbon atoms. Each of the hydrocarbons present in crude oil has a slightly different use. Mixed together they are of no use at all. It is necessary, therefore, to separate them before they can be used productively. Crude oil is separated into its different components by a process called fractional distillation. The products of fractional distillation are often converted into other, even more useful hydrocarbons by a process called cracking. 2. Fractional distillation The different hydrocarbons in crude oil have different boiling points. This is because the chain length varies. The greater the number of carbon atoms in the chain, the longer the chain length. This results in more Van der Waals forces acting between the molecules and a greater intermolecular attraction. Thus more energy is needed to separate the molecules and the boiling point is higher. It is the difference in boiling points of the different hydrocarbons in crude oil which is used to separate them from each other. The crude oil is passed into a tall tower called a fractionating column. This is very hot near the base but much cooler near the top. When the crude oil is passed into the tower, near the bottom, most of the mixture boils and starts to rise up the tower. As they rise up the tower, they start to cool down and will gradually condense back into liquid form. They are then tapped off. The larger hydrocarbons, with higher boiling points, will condense first and be tapped off near the base of the column. The smaller hydrocarbons, with smaller boiling points, will condense later and be tapped off near the top of the column. Thus the separation is achieved. Not that the process involves breaking intermolecular forces only; the molecules themselves are unaffected by this process. This process does not actually separate the crude oil mixture into pure hydrocarbon components, but into mixtures called fractions. Fractions are mixtures of hydrocarbons with similar boiling points. In many cases these fractions can be used directly, but sometimes further separation is required into purer components. The following page shows a diagram of a typical fractionating column, and a table showing the most important fractions and their main uses:
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94 A fractionating column
Fractions from crude oil Name of fraction Liquefied petroleum gas Petrol or gasoline Naphtha Kerosine or paraffin Diesel or gas oil Mineral/lubricating oil Fuel oil Wax and grease Bitumen or tar Boiling range / oC Less than 25 Number of hydrocarbons 14 Uses Gas for camping/ cooking Fuel for cars etc Petrochemicals Plane fuel, petrochemicals lorry, central heating fuel
Lubrication, petrochemicals
Ship fuel, power stations Candles, grease, polish Road surfaces, roofing
Above 450
More than 50
The term petrochemical means that the compounds are converted into other chemicals for use as solvents, paints and various other things.
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95 3. Cracking Although all of the fractions produced from crude oil have their uses, some of the fractions are produced in greater quantities than needed, whilst others are not produced in sufficient quantities. The table below gives an example of the difference between the supply and demand of some important fractions: Supply and demand for fractions Fraction Liquefied petroleum gases Petrol and naphtha Kerosine Gas oil Fuel oil and bitumen Approximate supply/% 2 16 13 19 50 Approximate demand/% 4 27 8 23 38
This disparity can be corrected by breaking up some larger hydrocarbons in fuel oil into the smaller ones found in gas oil, or by breaking up some hydrocarbons in kerosene into the smaller ones found in petrol, naphtha or the liquefied petroleum gases. In other words the larger fractions (for which supply exceeds demand) can be broken up into smaller fractions (for which demand exceeds supply). The process by which this is carried out is called cracking. Cracking has the added advantage of producing other useful hydrocarbons not naturally present in crude oil, such as alkenes (widely used as petrochemicals), cycloalkanes and branched alkanes (widely used in motor fuels) and aromatic hydrocarbons (used as petrochemicals and as motor fuels). Thus cracking is important for two reasons: i) It converts low-demand fractions into higher demand fractions ii) It makes useful hydrocarbons not naturally found in crude oil There are two types of cracking: thermal cracking and catalytic cracking. Both involve the breaking of C-C bonds to form smaller molecules. C-C bonds are weaker than C-H bonds and so break more easily when heated.
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96 a) Thermal cracking In thermal cracking, the bonds are broken using a high temperature (400 900oC) and a high pressure (70 atmospheres). The high temperatures mean that the molecule breaks near the end of the chain, giving a high percentage of small alkenes such as ethene. Most thermal cracking reactions involve the formation of one of more small alkane molecules and one alkene molecule. Naphtha (C7 C14) is usually used as the starting material. Eg C8H18 C6H14 + C2H4
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97 b) Catalytic cracking In catalytic cracking, the bonds are broken using a high temperature (450 oC, which is generally lower than in thermal cracking), a slight pressure (slightly greater than 1 atmosphere), and a zeolite catalyst. Catalytic cracking is cheaper and more efficient than thermal cracking as it uses a lower temperature and pressure. The zeolite catalyst favours the formation of branched alkanes and cycloalkanes, which are widely used in motor fuels. The most important product of catalytic cracking is 2-methylheptane, which is the major component of petrol. It also produces aromatic hydrocarbons such as benzene, which have a variety of uses.
H H H H H H H C H H C C H C H C C H H
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Eg
A table summarising the differences between thermal and catalytic cracking can is shown below: Type of cracking Conditions Thermal High temperature (400 900 oC) High pressure (70 atm) Catalytic High temperature (450 oC) Slight pressure ( > 1 atm) Zeolite catalyst Main products High percentage of alkenes Motor fuels (ie branched alkanes) Aromatic hydrocarbons
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A fuel is a something that can be changed in a reacting vessel to produce useful energy. Hydrocarbons, and especially alkanes, will react with oxygen in the air to give carbon dioxide and water. A reaction with oxygen is known as combustion. As alkanes are unreactive the reaction needs heat or a spark to get going. These reactions are very exothermic, which means that heat energy is released. This heat energy can be used for direct heating (eg camping gas, central heating, candles). It can also be converted into mechanical energy (eg cars, lorries, ships), or even electrical energy (eg power stations). Typical examples of combustion reactions include: Reaction CH4 + 2O2 CO2 + 2H2O C4H10 + 6O2 4CO2 + 5H2O C8H18+ 12O2 8CO2 + 9H2O Enthalpy change/ kJmol-1 -890 -2877 -5470
The release of heat energy during these combustion reactions results in their widespread use as fuels.
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The less oxygen that is available, the more likely it is that incomplete combustion will occur. This is a particular problem in internal combustion engines where the air supply is limited. Incomplete combustion is a problem for three reasons: i) ii) Less energy is released by incomplete combustion than by complete combustion. Carbon monoxide is a pollutant it is absorbed by the blood in place of oxygen, and hence reduces the ability of the blood to carry oxygen causing suffocation and eventually death. Carbon particles can cause breathing difficulties and cancer.
iii)
It is therefore desirable to ensure that the air supply is as good as possible when burning hydrocarbon fuels.
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100 Occasionally incomplete combustion is desirable such as with a Bunsen burner. Closing the air hole produces a yellow flame (the yellow colour results from hot carbon particles) and this makes the flame more visible and causes a more gentle heat. Usually, however, complete combustion is considered more desirable. d) Sulphur dioxide Most crude oil deposits contain sulphur as an impurity. Oil refineries are increasingly treating the petrol fractions to lower the sulphur content, but some sulphur is still present in most hydrocarbon fuels. When the fuel is burned, the sulphur also burns, producing sulphur dioxide: S(s) + O2(g) SO2(g) This gas dissolves in rainwater forming a very acidic solution, known as acid rain. This causes various problems, including erosion of buildings and statues, killing of plants and trees, and killing of fish through contamination of lakes. e) Oxides of nitrogen Most fuels are not burned in pure oxygen but in air, which contains 80% nitrogen. Although nitrogen is not a reactive gas, the high temperatures and the spark in combustion engines cause some of the nitrogen to react with the oxygen to produce nitric oxide and nitrogen dioxide: N2(g) + O2(g) 2NO(g) 2NO(g) + O2(g) 2NO2(g) Nitrogen dioxide (NO2) also dissolves in rainwater to form an acidic solution and contributes to the problem of acid rain. Nitrogen oxides can also combine with unburned hydrocarbons to produce a photochemical smog. f) Unburned hydrocarbons Some of the hydrocarbon fuel is vaporised in the engine but escapes before it is burned. These unburned hydrocarbons cause various problems. They are toxic and can cause cancer if breathed in. They also combine with oxides of nitrogen to produce a photochemical smog.
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Many factory chimneys contain alkaline materials such as lime (calcium oxide). These absorb the acidic gases such as SO2 and thus prevent them from escaping: SO2 + CaO CaSO3 Further reactions result in the formation of CaSO4 (gypsum) which is used to make plaster. b) Catalytic Converters
Most modern car exhausts are now fitted with catalytic converters. These are designed to convert some of the more harmful gases present in car exhausts into less harmful ones. Unburned hydrocarbons, carbon monoxide and the oxides of nitrogen can all be converted into less harmful gases inside these converters. There are two main types of reaction taking place in a catalytic converter: i) Removal of carbon monoxide and nitrogen monoxide 2NO(g) + 2CO(g) N2(g) + 2CO2(g) Hence harmful NO and CO gases are converted into the less harmful nitrogen and carbon dioxide. ii) Removal of unburned hydrocarbons and nitrogen monoxide eg. C8H18 + 25NO 8CO2 + 9H2O + 12.5N2 Hence harmful unburned hydrocarbons and oxides of nitrogen are converted into the less harmful carbon dioxide, water and nitrogen.
The reality is, however, that the burning of hydrocarbon fuels has caused and continues to worsen most of the planets most serious environmental problems. Although technological innovations such as catalytic converters can limit some of the damage, the only action which will have any lasting effect is to reduce the reliance of rich Western countries, especially the USA, on fossil fuels. This will only happen if the potential of alternative sources of energy is more fully exploited, the political and economic power of oil barons is curbed and wealthy industrialised countries look at ways to reduce their energy consumption. Achieving these goals, however, has been socially and politically problematic.
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