Edexcel Chemistry AS Notes

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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 1

1.1 Atomic structure

The Structure of the Atom

Mass Spectrometry

Electronic Structure

Ionisation Energies

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THE STRUCTURE OF THE ATOM

a) Protons, neutrons and electrons

Atoms are made up of three fundamental particles: protons, neutrons and

electrons.

Protons and neutrons are found in the nucleus and are collectively called

nucleons. Electrons orbit the nucleus in a similar way to that in which planets orbit

a sun. In between the electrons and nucleus there is nothing (empty space).

The nucleus is very small; if an atom were the size of a football pitch, the nucleus would be the size of a drawing pin.

The basic properties of these three particles can be summarized in the following table:

Particle Charge Mass

Proton +1 unit Approx 1 unit

Neutron No charge Approx 1 unit

Electron -1 unit Approx 1/1840 units (very small)

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.

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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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: 12 protons, 12 neutrons 25 12Mg: 12 protons, 13 neutrons 26 12Mg: 12 protons, 14 neutrons

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 12Mg2+: 12p, 12n, 10e 24 12Mg+: 12p, 12n, 11e 24 12Mg 12p, 12n, 12e 24 12Mg-: 12p, 12n, 13e

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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 Mass number Relative isotopic mass 1 1H 1 1.006 4 2He 4 4.003 9 4Be 9 9.012 27 13Al 27 26.919 59 27Co 59 58.933

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 Common mass numbers Relative atomic mass

Mg 24, 25, 26 24.32

Cl 35, 37 35.45

Br 79, 81 79.91

Ba 134, 135, 136, 137, 138 137.33

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.

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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 depends on m/e ratio (DEFLECTION)

The heavier the particle, the less the deflection

5- Electric current measured as ions land on plate (DETECTION)

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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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 18 20 22 24 26 20 40 60 80 100 M/Z relative abundance

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2. Calculating relative atomic masses

The relative atomic mass can be calculated by the formula:

Σ (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 C5H12+, 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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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=1

n=2

n=3

n=4

n=6

n=5

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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s-orbitals: these are spherical.

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, s-orbitals have the lowest energy and the energy of the other s-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 sub-level 2nd sub-level 3rd sub-level 4th sub-level 5th sub-level 1 1 x 1s 2 1 x 2s 3 x 2p 3 1 x 3s 3 x 3p 5 x 3d 4 1 x 4s 3 x 4p 5 x 4d 7 x 4f 5 1 x 5s 3 x 5p 5 x 5d 7 x 5f 9 x 5g

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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 levels of some energy levels soon overlap with the low energy sub-levels of higher energy sub-levels, resulting in a more complex energy level diagram:

1s 2s 3s 4s 2p 3p 3d 4p 4d 4f n=1 n=2 n=3 n=4 E N E R G Y

Starting with the lowest energy, the orbitals can thus be arranged as follows:

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d

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.

1s│2s 2p│3s 3p │ 4s 3d 4p │5s 4d 5p│6s 4f 5d 6p

The arrangement of shells and the maximum number of electrons in each can be summarised as follows:

Shell number Orbitals in shell

1 1 x1s 2 1 x 2s, 3 x 2p 3 1 x 3s, 3 x 3p 4 1 x 4s, 5 x 3d, 3 x 4p 5 1 x 5s, 5 x 4d, 3 x 5p 6 1 x 6s, 7 x 4f, 5 x 5d, 3 x 6p

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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 Number of electrons in each sub-level Max. no of electrons

1 2 x 1s 2 2 2 x 2s, 6 x 2p 8 3 2 x 3s, 6 x 3p 8 4 2 x 4s, 10 x 3d, 6 x 4p 18 5 2 x 5s, 10 x 4d, 6 x 5p 18 6 2 x 6s, 14 x 4f, 10 x 5d, 6 x 6p 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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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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: 1s1 He: 1s2 Li: 1s22s1 Be: 1s22s2 B: 1s22s22p1 C: 1s22s22p2 or 1s22s22p63s23px13py1 N: 1s22s22p3 or 1s22s22p63s23px13py13pz1 O: 1s22s22p4 or 1s22s22p63s23p23px23py13pz1 F: 1s22s22p5 Ne: 1s22s22p6 Na: 1s22s22p63s1 Mg: 1s22s22p63s2 Al: 1s22s22p63s23p1 Si: 1s22s22p63s23p2 or 1s22s22p63s23px13py1 P: 1s22s22p63s23p3 or 1s22s22p63s23px13py13pz1 S: 1s22s22p63s23p4 or 1s22s22p63s23px23py13pz1 Cl: 1s22s22p63s23p5 Ar: 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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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 14 4s 3d 4p K [Ar] ↑ Ca [Ar] ↑↓ Sc [Ar] ↑↓ ↑ Ti [Ar] ↑↓ ↑ ↑ V [Ar] ↑↓ ↑ ↑ ↑ Cr [Ar] ↑ ↑ ↑ ↑ ↑ ↑ Mn [Ar] ↑↓ ↑ ↑ ↑ ↑ ↑ Fe [Ar] ↑↓ ↑↓ ↑ ↑ ↑ ↑ Co [Ar] ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ Ni [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ Cu [Ar] ↑ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ Zn [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ Ga [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ Ge [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ As [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ Se [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ Br [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ Kr [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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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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:

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.

Variation of first ionisation energy with atomic number for the first twenty elements

0 500 1000 1500 2000 2500 0 5 10 15 20 atomic number fi rs t ion is a ti on e ne rgy ( k J pe r m ol e )

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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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.

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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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.

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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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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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) Successive ionisation energies

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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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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:

Variation of ionisation energy with number of ionisations for aluminium

100 1000 10000 100000 1000000 1 2 3 4 5 6 7 8 9 10 11 12 13 number of ionisations ion is a ti on e ne rgy

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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 22 Note that the largest jumps by far occur between the 3rd and 4th 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:

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.

Variation of ionisation energy with number of ionisations in silicon 100 1000 10000 100000 1000000 1 3 5 7 9 11 13 number of ionisations ion is a ti on e ne rgy

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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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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 Avogadro’s number. The number of moles can be calculated by dividing the number of particles by Avogadro’s number.

(Number of particles) = (number of moles) x L

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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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 25 The symbol for molar mass of compounds or molecular elements is mr. 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:

MASS = MOLAR MASS X NUMBER OF MOLES

or m = m

r

x n

Mass must be measured in grams and molar mass in gmol-1.

mass

moles molar

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REACTING MASSES

It is possible to use the relationship moles = mass/mr 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

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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 x 100 maximum amount of product possible

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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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 = 52.3 %

(2 x 40) + (1 x 2) + (1 x 71)

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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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 dm3 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:

Number of moles = volume x molarity

n = C x V

moles

volume(dm3₎ molarity

The volume of solution in this case must always be measured in dm3 (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.

Concentration (gdm

-3

) = Molarity x molar mass

Or C

g

= C

m

x m

r

C(grams) molarity

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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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 cm3 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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GASES

The volume occupied by a gas depends on a number of factors:

i) the temperature: the hotter the gas, the faster the particles are moving

and the more space they will occupy

ii) the pressure: the higher the pressure, the more compressed the gas will

be and the less space it will occupy

iii) 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 mr = 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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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:

use the ratios in the equation to find the number of moles of other species n = PV RT n = mass RMM n = CV n = particles L for solutions:

These relationships are frequently used in practical chemistry. Typical calculations in AS Practical Chemistry involve:

i) Determining the concentration of a solution

ii) Determining the relative molecular mass of a solid iii) 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 x 100 mass of impure substance

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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 : 14.2 12 1 = 7.15 : 14.2 7.15 : 7.15 = 1 : 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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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

Charge Formula Name

+1 Na+ Sodium +1 K+ Potassium +1 Ag+ Silver +1 H+ Hydrogen +1 NH4+ Ammonium +1 Cu+ Copper(I) +2 Mg2+ Magnesium +2 Ca2+ Calcium +2 Fe2+ Iron(II) +2 Zn2+ Zinc +2 Pb2+ Lead(II) +2 Cu2+ Copper(II) +2 Ni2+ Nickel(II) +3 Al3+ Aluminium +3 Cr3+ Chromium(III) +3 Fe3+ Iron(III)

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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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 35 ii) anions

Charge Formula Name

-1 OH- Hydroxide -2 SO42- Sulphate -2 CO32- Carbonate -1 NO3- Nitrate -1 HCO3- 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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ii) the reaction coefficients

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 A-level chemistry:

(s) - solid (l) - liquid (g) - gas

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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)

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Topic 1.3

BONDING

Types of bond

States of matter

Structure and physical properties

Molecular shapes

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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.

Na x Cl oo o o oo o Na + Cl oo o o oo o x -Na + Cl oo o o oo o x

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-Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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 H O H x o x_o x x _xx H O H

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 xx xo x o x o N H H H H + N H H H 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

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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.

Mg Mg 2+ 2+ e e e e

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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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 He 2.1 Li Be B C N O F Ne 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Na Mg Al Si P S Cl Ar 0.9 1.2 1.5 1.8 2.1 2.5 3.0 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 0.8 1.0 1.3 1.5 1.6 1.6 1.5 1.8 1.8 1.8 1.9 1.6 1.6 1.8 2.0 2.4 2.8

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 B

o

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 H

o

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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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)

O x_o H

- +

a polar covalent bond

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 xCl 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) Mg Mg x o x o 2+ 2+ 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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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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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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.

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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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.

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.

LIQUIDS

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IONIC STRUCTURES

Bonding in ionic compounds

An ionic bond is an attraction between oppositely charged ions. After the ions are formed they all come together to form a lattice. A lattice is an infinite and repeating arrangement of particles. All the anions are surrounded by cations and all the cations are surrounded by anions.

The coordination number of an ion in an ionic solid is the number of oppositely charged ions which surround it. This varies from substance to substance but is usually 4, 6 or 8.

Example – sodium chloride

In sodium chloride, NaCl, each sodium ion is surrounded by six chloride ions and vice versa.

The diagram below shows the structure of sodium chloride. The pattern repeats in this way and the structure extends (repeats itself) in all directions over countless ions. You must remember that this diagram represents only a tiny part of the whole sodium chloride crystal.

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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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 48 Na+ Cl -Na+ Na+ Na+ Na+ Na+ Cl -Cl -Cl -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 NaCl MgO

Melting point/oC 801 2852 Boiling point/oC 1459 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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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 Mg2+ Mg2+ Mg2+ Mg 2+ Mg2+ Mg 2+ Mg2+ _Mg2+ Mg2+ e e e e e e e e e e e e e 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 Na K Be Mg Melting point/ oC 98 64 127 8 649 Boiling point/ oC 883 760 297 0 110 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

This is a simplified 2D form of the metal lattice

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Kendhikulhudhoo School-AS Chemistry Notes-2010 Page 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.