Preview Cambridge Pre-U Chemistry Coursebook

Cambridge Pre-U

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Chemistry

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Coursebook

Cambridge Elevate edition Original material Š Cambridge University Press 2016

Cambridge Pre-U Chemistry

S3: Electrons in atoms understand and use the relationship between group number and successive ionisation energies

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Learning Outcomes

S3.1 Simple electronic structure

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We have seen that atoms consist of a nucleus (containing protons and neutrons) and electrons which surround the nucleus at certain distances away on shells. These shells represent energy levels and as a result the electron energy is quantised. A quantum is a packet of energy implying that only discrete energy levels are possible. As we will see in chapter S37 on electronic spectra, electrons can only ‘move’ between certain energy levels.

In other words for an electron to occupy a higher energy a distinct amount of energy must be supplied or is given out. We can assign each energy level, or shells, a quantum number, the principal quantum number, symbol n, which is an integer with n ≥ 1. The shells increase in radius with increasing n and thus move further away from the nucleus, which weakens the attraction between the positive nucleus and the negative electrons in the shells. Each shell can take a maximum number of electrons: the first shell 2, the second shell 8, the third shell 18, the fourth 32, etc.

Shell number, n

Maximum number of electrons

1

2

2

18

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3

8

4

32

n

2n2

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Table S3.1 Maximum numbers of electrons for the first five electron shells.

The arrangement of electrons in an atom is called its electronic structure or electronic configuration. The electronic configurations of lithium, carbon and neon are shown in Figure S3.1, together with a shorthand way of writing this structure.

Figure S3.1 The simple electronic structures of lithium, neon and chlorine. The nuclei of the atoms are not shown.

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Therefore, we can write electronic configurations for atoms as follows: second shell first shell e.g. chlorine, Cl (17e–): 2,

third shell 8,

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QUESTIONS

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3.2 Write the electronic configuration in box notation for the following atoms: C, Na, S, K

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3.1 Write the simple electronic configuration of the following atoms: H, He, B, C, Ne, Na, Al, S, Ar

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This is called box notation in which electrons are assigned to their principal shell. Later we will use a more detailed box notation. A simple electronic configuration can be written as: 2,8,7 for the chlorine atom.

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Ions can be written in a similar fashion: for example, the chloride ion: Cl− (18e−): 2,8,8 QUESTION

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3.3 Write the simple electronic configuration of the following ions: H+, H−, Li+, C4+, O2−, Na+, Na−, Al3+, S2−

We can see that the hydride ion (H−), the lithium ion, Li+, and the carbon ion, C4+, have the same electronic configuration (‘2’); we call them isoelectronic. The same applies to oxide, O2−, and aluminium ion, Al3+.

QUESTION

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3.4 Predict an ion that is isoelectronic with the nitrogen atom.

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Table S3.2 shows the number of electrons in each of the principal quantum shells (energy levels) for the first 11 elements in the Periodic Table. Each principal quantum shell can hold a maximum number of electrons: • shell 1 – up to 2 electrons • shell 2 – up to 8 electrons • shell 3 – up to 18 electrons • shell 4 – up to 32 electrons.

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Atomic number

Number of electrons in shell n=1

n=2

n=3

1

1

0

0

He

2

2

0

0

Li

3

2

1

0

Be

4

2

2

0

B

5

2

3

0

C

6

2

4

0

N

7

2

5

0

O

8

2

6

F

9

2

7

Ne

10

2

8

Na

11

2

8

N

H

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Element

0 0

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0 1

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Table S3.2 Simple electronic configurations of the first 11 elements in the Periodic Table.

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S3: Electrons in atoms write equations for successive ionisation energies

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Learning Outcomes

S3.2 Evidence for electronic structure

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Ionisation energy, ΔHIE

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In the Coursebook, the concept of successive ionisation energies is introduced. In this section, we revisit this idea using the example of magnesium.

Scientists can work out how much energy has to be supplied to form an ion by bombarding the atom with high energy electrons. The electron impact will remove an electron from the atom and form an ion. The energy change that accompanies this process is called the ionisation energy.

The first ionisation energy of an element is the energy needed to remove one mole of electrons from one mole of atoms each in the gaseous state to form one mole of gaseous 1+ ions.

The first ionisation energy can be expressed as ΔHIE1. Using calcium as an example: First ionisation energy: Ca(g) → Ca (g) + e +

−

ΔHIE1 = 590 kJ mol−1

If a second electron is removed from each ion in a mole of gaseous 1+ ions, we call it the 2nd ionisation energy, ΔHi2.

Again, using calcium as an example:

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Second ionisation energy: Ca+(g) → Ca2+(g) + e− ΔHIE2 = 1150 kJ mol−1

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Removal of a third electron from each ion in a mole of gaseous 2+ ions is called the 3rd ionisation energy. Again, using calcium as an example: Third ionisation energy: Ca2+(g) → Ca 3+(g) + e− ΔHIE3 = 4940 kJ mol−1

We can continue to remove electrons from an atom until only the nucleus is left. We call this sequence of ionisation energies, successive ionisation energies. The successive ionisation energies for the first 12 elements in the Periodic Table are shown in Table S3.3 and a plot of the 11 ionisation energies for magnesium is shown in Figure S3.2.

The data in Table S3.3 show us that: • For each element, the successive ionisation energies increase.

• There is a big difference (jump) between some successive ionisation energies. For nitrogen this jump occurs between the fift h and sixth ionisation energies. For magnesium the first big difference occurs between the second and third ionisation energies. The jumps can be seen as an indirect proof for the shell structure of the atom.

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Ionisation energies in kJ mol−1 for number of electrons removed Element

1

2

3

4

5

6

7

8

9

10

1

H

1310

2

He

2370

3

Li

519

7300 11 800

4

Be

900

1760 14 850 21 000

5

B

799

2420

3660 25 000 32 800

6

C

1090

2350

4620

6220 37 800 47 300

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N

1400

2860

4580

7480

8

O

1310

3390

5320

7450 11 000 13 300 71 300 84 100

9

F

1680

3370

6040

8410 11 000 15 200 17 900 92 000 106 000

10

Ne

2080

3950

6150

9290 12 200 15 200 20 000 23 000 117 000 131 400

11

Na

494

4560

6940

9540 13 400 16 600 20 100 25 500

12

Mg

734

1451

7733 10 540 13 628 17 995 21 704 25 656

11

12

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5250

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N

9450 53 300 64 400

28 900 141 000 158 700 31 643

35 462 169 991 189 367

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Table S3.3 Successive ionisation energies for the first 12 elements in the Periodic Table.

Figure S3.2 The successive ionisation energies of magnesium.

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S3: Electrons in atoms ■

describe the concept of shielding in multi-electron atoms use the relationship between group number and successive ionisation energies

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Learning Outcomes

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S3.3 Three factors that influence the successive ionisation energies of an atom The size of the nuclear charge

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As the atomic number (number of protons) increases, the positive nuclear charge increases.

The bigger the positive charge, the greater the attractive electrostatic force between the nucleus and the electrons. So, more energy is needed to overcome these attractive forces if an electron is to be removed.

In general, ionisation energy increases as the proton number increases.

Distance of outer electrons from the nucleus

The force of attraction between positive and negative charges decreases rapidly as the distance between them increases. So, electrons in shells further away from the nucleus are less attracted to the nucleus than those closer to the nucleus.

The further the outer electron shell is from the nucleus, the lower the ionisation energy.

This is related to the shielding effect, discussed below.

Shielding effect and repulsion of electrons

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As all electrons are negatively charged, they repel each other. Electrons in (full) inner shells, the core electrons, repel electrons in shells further away from the nucleus; they prevent the full nuclear charge being felt by the outer electrons. This effect is often called shielding (Figure S3.3). Shielding can be understood as electrons in subshells below an (valence) electron partially cancelling the positive charge of the nucleus.

The ionisation energy reduces as the number of full electron shells (core electrons) between the outer electrons and the nucleus increases.

Figure S3.3 The shielding of lower (sub)shells to the valence electron in a higher energy level in lithium.

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Shielding should not be understood as an actual ‘shield’ between the core electrons and the valence electrons, it is more like the sum of the number of positive protons in the nucleus and the number of electrons between the nucleus and the electrons in question. The core electrons cancel some of the force of attraction between the nucleus and the valence electrons.

For example, in lithium the two electrons in the first shell will shield the valence electron in the second shell by almost one unit charge each. As a result, it is a good approximation to assume that the valence electron in lithium experiences an attractive electrostatic force of +1 rather than +3; we sometimes call this the effective nuclear charge, Zeff. A proper treatment of effective nuclear charges would require us to use Slater’s rules, which go beyond the scope of this book. Please note that electrons on the same energy level the same nuclear charge, so the two valence electrons in beryllium experience roughly the same +2 charge (a little less due to the repulsion between them).

While electrons in different shells repel each other, the same is also true for electrons in the same shell. We will look at the He+ ion to understand how this occurs. This ion has one electron in the valence shell just like hydrogen (we say that it is isoelectronic to hydrogen, or hydrogenic). This electron is attracted to two protons in the nucleus and therefore possesses a certain potential energy, which is associated with a certain shell radius, r1 (Figure S3.4).

Figure S3.4 The effect of an additional electron on the shell radius (not to scale): r1 < r2.

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Let us add another electron to form the helium atom. The attractive electrostatic (Coulomb) force of the nucleus is the same for both electrons but due to the repulsion between the two electrons the shell radius, r2, increases (we can interpret this as the two electrons trying to get out of each other’s way). You could say that the atom gets slightly inflated: r1 > r2. As a result, it is easier to ionise the first electron in the helium atom (either of the two electrons present), as it is further away from the nucleus than a single electron would be on its own. Once one electron has been removed, the second electron experiences no repulsion, so it moves closer to the nucleus, which causes an increase in the second ionisation energy.

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QUESTIONS Table S3.4 outlines the ionisation energies (kJ molâ&#x2C6;&#x2019;1) of five elements.

1st ionisation energy

2nd ionisation energy

3rd ionisation energy

4th ionisation energy

A

520

7301

11 817

-------

B

578

1817

2746

10 813

C

1087

2354

4621

6425

D

496

4566

6917

9547

E

590

1146

4944

6469

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3.5

Table S3.4 Ionisation energies of five elements.

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a Which of these five elements, when it reacts, is most likely to form a 3+ ion? Explain your answer. b Which of these five elements are likely to be in the same group of the Periodic Table?

Explain your answer and state the group.

Suggest the possible group(s) of the Periodic Table in which element C could be found, and explain your answer briefly.

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With these points in mind we can now interpret the ionisation energies of atoms.

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S3: Electrons in atoms understand and use the relationship between group number and successive ionisation energies

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Learning Outcomes

S3.4 Interpreting successive ionisation energies

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Figure S3.5 shows a graph of successive ionisation energies against the number of electrons removed for sodium. A logarithmic scale (to the base 10) is used because the values of successive ionisation energies have such a large range.

Figure S3.5 Graph of logarithm (log10) of ionisation energy of magnesium against the number of electrons removed.

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We can deduce the following about magnesium from Figure S3.5: • The first electron removed has the lowest 1st ionisation energy, when compared with the rest of the data. It is very easily removed from the atom. It is therefore likely to be a long way from the nucleus and well shielded by inner electron shells and also shielded by the second electron in the same shell.

• The second electron is only slightly harder to remove as it is lacking the repulsion from the first electron.

• There is a big jump in the value of the ionisation energy after the second electron. This suggests that the third electron is in a shell closer to the nucleus than the second electron.

Taken together, the first three ionisation energies suggest that magnesium has two electrons in its outer shell.

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â&#x20AC;˘ From the third to the 10th electrons removed, there is only a gradual change in successive ionisation energies. This suggests that all these eight electrons are in the same shell. â&#x20AC;˘ The 11th and 12th electrons have extremely high ionisation energies, when compared with the rest of the data. This suggests that they are very close to the nucleus. There must be a very great force of attraction between the nucleus and these electrons and there are no inner electrons to shield them. The large increase in ionisation energy between the 10th and 11th electrons confirms that the 11th electron is in a shell closer to the nucleus than the 10th electron.

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Figure S3.6 shows this arrangement of electrons.

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Figure S3.6 The arrangement of electrons in an atom of magnesium can be deduced from the values of successive ionisation energies.

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S3: Electrons in atoms recall the relationship between shell number and number of subshells

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Learning Outcomes

S3.5 Subshells and atomic orbitals

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We saw in section S3.1 that the number of electrons each shell can take is determined by the expression 2n2. There is a reason for the formula 2n2 and therefore the maximum number of electrons each shell can take: quantum mechanics, the modern theory behind the structure of the atom and its electrons. The theory is based on the famous Schrödinger equation. From the Schrödinger equation it turns out that each of the shells may consist of subshells (abbreviated s, p, d, f, g, …). In the hydrogen atom this makes no difference, as all the subshells belonging to a principal quantum number are degenerate, that is, they have the same energy (see Figure S3.7). There is only one s subshell, three p subshells, five d subshells, seven f subshells, etc. per principal quantum number. As we can see, the number of subshells increases with its principal quantum number (for example the first shell only has the s subshell, the second shell has one s and three p subshells, etc.). n

3

d p s

2

p s

1

s

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energy

4

f d p s

subshells

Figure S3.7 The subshells for the first four principal quantum numbers.

The labels s, p, d and f arise from the line spectra of atoms and stand for sharp, principal, diff use and fundamental (see also Chapter S37). We will see later that even more subshells are possible provided there are enough electrons to occupy them. As already discussed the first principal quantum level, n = 1, can hold a maximum of 2 electrons in an s subshell. The second principal quantum level, n = 2, can hold a maximum of 8 electrons: 2 electrons in the s subshell and 6 electrons in the p subshell. The third principal quantum level, n = 3, can hold a maximum of 18 electrons: 2 electrons in the s subshell, 6 electrons in the p subshell and 10 electrons in the d subshell and so on.

With the introduction of subshells, we can write more detailed electron configurations.

For the chlorine atom we used to write: Cl (17e−): 2, 8, 7. Using subshells this can be expanded to: Cl (17e−): 1s2 2s2 2p6 3s2 3p5 showing the occupancy of each of the subshells.

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QUESTIONS 3.6 Write the electronic configuration for the following atoms: H, He, B, C, Ne, Na, Al. 3.7 Write the electronic configuration for the atoms with the atomic number 16, 9 and 18.

3.8 Write the electronic configuration in box notation for the following atoms: C, Na, S, K. Ions can be written in a similar fashion: for example the chloride ion: Cl− (18e−): 1s2 2s2 2p6 3s2 3p6.

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3.9 Write the electronic configuration of the following ions:

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H−, Li+, C4+, O2−, Na+, Na−, Al3+, S2−

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S35: Stereochemistry

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understand that chiral molecules rotate the plane of polarised light explain the +/− notation of chiral molecules use the Cahn–Ingold–Prelog priority rules to assign R/S to chiral centres define the terms enantiomer, diastereomer and meso compound in the case of molecules with two chiral centres.

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Learning Outcomes

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S35.1 Introduction

Figure S35.1 Our hands, like some molecules, are chiral objects.

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Stereochemistry is concerned with the chemistry of stereoisomers – molecules with the same atoms bonded to each other but with different arrangements of the atoms in space (see S14.1, 14.2, the section on ‘Stereoisomerism’). In this chapter you will find out more about optical isomers.

The human hand is a familiar chiral object – the right hand is a non-superimposable mirror image of the left hand. The word chirality is in fact derived from the Greek word for hand. Most amino acids, from which our proteins are constructed, also have the property of being nonsuperimposable with their mirror image and are also chiral, or are said to possess molecular ‘handedness’. This is important because, just as a left-handed glove will only fit a left hand, a receptor in a biological system will require a particular molecular mirror image for a correct fit.

Therefore, chemists must be concerned about which isomer they should synthesise.

S35.2 Optical activity

We have already seen that if a molecule contains a carbon atom that is bonded to four different atoms or groups of atoms, it can form two isomers. The two different molecules have no planes of symmetry and are non-superimposable mirror images. The carbon atom with the four different groups attached is called the chiral centre of the molecule (see S14.1, 14.2, the section on ‘Optical isomerism’). These two isomers are called enantiomers or optical isomers, because they are able to rotate the plane of polarised light, by the same amount to the left or to the right (see Figure S35.2).

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polarising filter angle of rotation of plane of polarization

unpolarised light

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rotated polarised light ght

light source polarised light

optically active solution

polariser axis

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Figure S35.2 The effect of an optically active solution on the plane of polarisation of planepolarised light. Plane-polarised light is passed through a tube containing a solution in an instrument called a polarimeter. If the compound is optical active, the plane of the light will be rotated on its way through the tube.

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Enantiomers which rotate the plane of polarised light in a clockwise direction have a positive (+) optical rotation. In contrast, enantiomers which rotate polarised light in an anticlockwise direction have a negative (â&#x2C6;&#x2019;) optical rotation (Figure S35.3). Other physical properties of the enantiomer pair will be identical. A mixture containing equal amounts of the two enantiomers is known as a racemic mixture. A racemic mixture has no effect on planepolarised light.

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Figure S35.3 Optical isomers of the amino acid alanine. The pair of enantiomers will rotate planepolarised light to the same extent in opposite directions.

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The property of optical activity was first observed by physicist Jean-Baptiste Bidot and was further investigated by Louis Pasteur. In 1848, Pasteur made a series of observations which would lead him to propose that molecular structure was responsible for the optical activity (Figure S35.4), laying the foundations of modern stereochemistry.

Figure S35.4 Louis Pasteur separated mirror image crystals of a tartaric acid salt from each other and showed that they rotated the plane of plane-polarised light to the same extent in opposite directions.

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Using optical activity

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The optical rotation observed for a solution of a particular chiral compound depends on the distance the light travels through the sample and the concentration of the enantiomer.

In order to compare measured rotations, a quantity called the specific rotation [α], can be calculated using the observed optical rotation (α), the concentration of the sample (c, g cm−3) and the path length of the light (l, dm). [α] = α l ×c

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The specific rotation is also dependent on the solvent, the temperature and the wavelength of light employed. Typically, optical rotations are measured at 20 °C in a solvent such as ethanol. The light used is from a sodium lamp which has a wavelength of 589 nm.

Comparing the specific rotation of a sample to that of the pure enantiomer can provide a means of determining its optical purity or enantiomeric excess, i.e. the extent to which one enantiomer is in excess in the sample. The maximum optical rotation will be observed with the single enantiomer, while a racemic mixture will have no optical rotation. The optical rotation of a non-racemic mixture of two enantiomers will depend on which enantiomer dominates the mixture and to what extent: % optical purity = observed [α] maximum [α] × 100

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Since chemical and biochemical reactions can result in changes in optical rotation, these reactions can be monitored using a polarimeter to measure optical rotation. For example, sucrose can undergo acid catalysed hydrolysis to glucose and fructose:

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Since optical rotation is directly proportional to concentration and the starting sucrose has a different optical rotation in comparison to the glucose/fructose product mixture, the change in optical rotation with time can be used to monitor this reaction.

S35.3 Assigning chiral centres

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Just as we can name alkene stereoisomers by inspecting their structures (see Chapter 14, section on ‘Cis–trans isomerism’), it is useful to be able to distinguish between enantiomers without taking optical rotation measurements. In 1966, a method of unambiguously assigning the ‘handedness’ of molecules was developed by three chemists, R.S. Cahn, C. Ingold and V. Prelog. This method labels each chiral centre R or S according to Cahn– Ingold–Prelog (CIP) priority rules, a system by which substituents are each assigned a priority.

Step 1:

All groups attached to the chiral carbon are assigned a priority. An atom with a higher atomic number has a higher priority than an atom with a lower atomic number. Hydrogen is the lowest possible priority substituent, because it has the lowest atomic number.

If two atoms attached to the chiral carbon are the identical, proceed along the two chains until there is a point of difference. The chain which has the first connection to an atom with the highest atomic number has the higher priority.

When isotopes are present, the atom with the higher atomic mass receives higher priority.

Step 2: When the groups attached to the chiral carbon have been assigned priorities 1–4, the molecule is orientated so that the lowest priority group, the group with priority 4, points away from the observer, i.e. it becomes the dashed line.

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Step 3:

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Determine the direction of an arrow that proceeds from groups 1 to 2 to 3. If this motion is clockwise, the chiral carbon is labelled R. If the arrow from 1 to 2 to 3 is anticlockwise, the chiral carbon has S configuration.

It is important to remember that there is no connection between the R or S assignment of a chiral centre and the experimentally measured optical rotation of the compound. There is no way of knowing whether a molecule will rotate the plane of polarised light clockwise or anticlockwise, and to what extent, by looking at the structure.

WORKED EXAMPLE S35.1

CH3

CH3

H

COOH

(–)-alanine

Step 1: Assign priorities

(+)-alanine

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N has the highest atomic number, so the NH2 group is given priority number 1.

3

3

CH3

1

2

H2N

COOH

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H

NH2

HOOC H

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H2N

The COOH group is given a higher priority than the CH3 group as the oxygen on the COOH carbon has a higher atomic number than the H on the CH3 carbon.

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mirror plane

Determine the configuration, R or S, of the enantiomers of alanine, as shown in the following figure.

4

CH3

2

1

NH2

HOOC H 4

4

Step 2: Orientate the molecule

Turn the molecule so that the group with priority number 4, in this case H, is pointing away.

3

3

CH3

2

4

H NH2

COOH

1

CH3 4

2

H H2N

COOH

1

R

Step 3: Determine the direction from 1 to 2 to 3 Clockwise = R

3

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CH3

Anticlockwise = S

2

HOOC

CH3 4

H NH2 1

(R)-(–)-alanine

4

H H2N

2

COOH

1

(S)-(+)-alanine

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S35.4 Molecules with two or more chiral centres

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Many organic molecules contain more than one chiral centre. For example, cholesterol has eight chiral carbons, marked * in Figure S35.5.

Figure S35.5 Cholesterol has eight chiral centres.

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In the case of 2,3-dichloropentane, two chiral centres are present. A pair of enantiomers of this compound can be seen in Figure S35.6. The steps previously described can be used to assign the chiral centres.

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Figure S35.6 2,3-dichloropentane has two chiral centres. The molecule can be redrawn to clearly show the configurations of both chiral carbons. Two enantiomers are shown.

To go from one enantiomer to the other, both chiral centres are inverted. However, what happens if only one chiral centre is inverted? The compounds produced are shown in Figure S35.7.

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Figure S35.7 Inverting just one chiral centre at a time produced two new compounds. They are not mirror images of the original (2R,3R)-2,3-dichloropentane.

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These two new compounds in Figure S35.7 are not mirror images of the original (2R,3R)2,3-dichloropentane. Stereoisomers that are not mirror images are called diastereoisomers.

Diastereoisomers have different physical and chemical properties, while the pair of enantiomers differ only in the direction in which they rotate polarised light. Figure S35.8 shows how the four possible stereoisomers of 2,3-dichloropentane are linked to one another.

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Figure S35.8 The four stereoisomers of 2,3-dichloropentane. There are two pairs of enantiomers.

The pairs of stereoisomers which are not enantiomers are diastereoisomers.

In the example above, a molecule with two chiral centres has four stereoisomers and two pairs of enantiomers. For a molecule with n chiral centres, the maximum number of stereoisomers can be calculated by simply calculating 2n. For example, a molecule with three chiral centres will have a maximum of 23 = 8 possible stereoisomers, four pairs of enantiomers: R,R,R

R,R,S

R,S,S

R,S,R

S,S,S S,S,R S,R,R S,R,S The structure and stereoisomers of tartaric acid are shown in Figure S35.9. Two chiral centres are present, so a maximum of four stereoisomers should be possible. However, two of the potential isomers are in fact the same molecule and although this molecule contains chiral centres, it is achiral because it is superimposable with its mirror image. R,S-tartaric acid becomes the S,R-tartaric acid when the molecule is rotated by 180o. Compounds which contain chiral centres but are themselves achiral are called meso compounds. This means

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that there must be a plane of symmetry in the molecule. Tartaric acid therefore has three stereoisomers – (R,R)-tartaric acid, (S,S)-tartaric acid and meso-tartaric acid.

Summary

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Chiral molecules have non-superimposable mirror images. The molecules are called enantiomers or optical isomers and they rotate the plane of plane-polarised light by the same amount in opposite directions.

Molecules which rotate the plane of polarised light in a clockwise direction are labelled +.

Molecules which rotate the plane of polarised light in an anticlockwise direction are labelled −.

Chiral centres can be assigned R or S by following Cahn–Ingold–Prelog (CIP) priority rules.

Changes in optical activity can be used to monitor reactions and test for chiral purity.

Stereoisomers that are not enantiomers are called diastereoisomers.

Compounds that contain chiral carbons but which are also achiral are called meso compounds.

The importance of chirality in drug design and methods of preparing pure enantiomers are covered in Chapter 30.

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Figure S35.9 The stereoisomers of tartaric acid. (R,S)-tartaric acid and (S,R)-tartaric acid are not enantiomers, they are identical and achiral. There are therefore three stereoisomers of tartaric acid – one pair of enantiomers and one meso compound.

■ ■ ■

■

Original material © Cambridge University Press 2016

7

Cambridge Pre-U Chemistry

End-of-chapter questions

LY

S35.1 For the following two compounds draw all possible stereoisomers, indicating those which are enantiomers and those which are diastereomers (use wedged and dashed bonds to distinguish between the individual isomers).

For each compound, assign R/S labels to the stereogenic centres. For which of these compounds does a meso isomer exist? a SH

N

OH

b

O

Br

FO R

R

EV IE

W

Br

Original material Š Cambridge University Press 2016

8

LY

Cambridge International Pre-University Chemistry

Syllabus Mapping Document Syllabus content

Present?

Location

Coursebook

(C)oursebook or (S)upplement or (B)oth

Chapter number

Supplement

Chapter title

Section number

Attention is drawn to a number of distinctive approaches within the syllabus, which are not commonly examined at this level. These topics should allow students to develop an important insight into chemical processes and should help to rationalise new and unfamiliar compounds and reactions.

EV IE

Physical chemistry

End page

W

• Van Arkel diagrams These diagrams put all binary compounds, be they ionic, covalent, giant covalent, metallic or semiconducting, on a single graph which explains their bonding type in terms of the difference in electronegativity and the average electronegativity of their constituent elements. It thus unifies all the chemical bonding types through an easily comprehensible concept.

• Functional group level This idea provides a simple framework through which a vast number of organic reactions can be understood: simply the number of bonds from a carbon atom to an electronegative element. Functional groups can be classified very simply by this scheme and so reactions between functional groups can thus immediately be understood as reduction, oxidation, or a reaction within a level such as substitution, hydrolysis or condensation.

• Orbitals An appreciation that electrons occupy orbitals in molecules as well as in atoms leads to the concept of bonding and anti-bonding molecular orbitals and an understanding that putting electrons into a bonding orbital leads to the nuclei being held together (i.e. bonded) whereas putting electrons into an anti-bonding orbital weakens bonding by pulling the nuclei apart from one another as these orbitals lie outside the internuclear region. An appreciation that vacant anti-bonding molecular orbitals exist is vital for a proper understanding of molecular electronic spectroscopy, and for understanding mechanism in further study.

1

Start page

O

Distinctive approaches

Section title

N

Section number

Part A of this section gives students insight into the structure of atoms and molecules and the forces between them, and the theory and measurement of enthalpy changes. In Part B, students are exposed to the driving forces of chemical reactions, the mechanisms of reactions and how their rate is studied, and the analysis of equilibrium conditions.

Throughout this study, students should be encouraged to consider the importance of evidence gathering and the limitations of theories.

Part A A1.1

A1.1 Introductory physical chemistry

students should be able to:

R

The following material may or may not have been covered as part of a Level 2 course (GCSE, IGCSE, O Level). If it has not previously been covered, it should be taught as part of this course and may be assessed.

(a) balance equations, including ionic equations, and use stoichiometric relationships

Y

C

1

Moles and equations

Reacting masses; The stoichiometry of a reaction

6

8

A1.1b

(b) use amount of substance (in moles) and Avogadro’s constant

Y

C

1

Moles and equations

Amount of substance

5

6

A1.1c

(c) calculate reacting masses from chemical formulae and equations

Y

C

1

Moles and equations

Amount of substance

5

6

A1.1d

(d) appreciate relative mass (atomic, isotopic, molecular and formula)

Y

C

1

Moles and equations

Masses of atoms and molecules

2

5

A1.1e

(e) appreciate relative charges, masses and location of protons, neutrons and electrons, and their behaviour in an electric field

Y

C

2

Atomic structure

Inside the atom

25

28

FO R

A1.1a

Original material © Cambridge University Press 2016

To add to which chapter of Coursebook?

Section number

Title

Coursebook Chapter number

Chapter title

A1.1f

(f) appreciate proton, electron, neutron and nucleon (mass) numbers in atoms and ions

Y

C

2

Atomic structure

A1.1g

(g) appreciate isotopes and their composition in terms of protons and neutrons

Y

C

2

Atomic structure

A1.1h

(h) appreciate electron sharing in covalent bonding, transfer in ionic bonding, and delocalisation in metallic bonding

Y

C

4

Chemical bonding

A1.1h

(h) appreciate electron sharing in covalent bonding, transfer in ionic bonding, and delocalisation in metallic bonding

Y

C

4

Chemical bonding

A1.1i

(i) understand the terms endothermic and exothermic and how they relate to bond making and bond breaking

Y

C

6

A1.1i

(i) understand the terms endothermic and exothermic and how they relate to bond making and bond breaking

Y

C

6

A1.1j

(j) appreciate the concepts of reversible reactions and dynamic equilibria

Y

C

8

A1.1k

(k) appreciate the kinetic-molecular model of change of state

Y

C

A1.1l

(l) understand simple collision theory as an explanation for factors affecting reaction rate.

Y

C

A1.2

A1.2 Atomic structure Content • atomic shells, subshells, orbitals and electron spin • aufbau principle and electron configurations • periodic trends in atomic properties

students should be able to:

R

• ionisation energies

Supplement

Section number

Section title

Start page

End page

To add to which chapter of Coursebook?

Section number

Title

N

Location (C)oursebook or (S)upplement or (B)oth

Numbers of nucleons

28

28

Isotopes; How many protons, neutrons and electrons?

28

29

Types of chemical bonding; Ionic bonding; Covalent bonding

49

55

Metallic bonding

58

60

Enthalpy changes

What are enthalpy changes?

90

91

Enthalpy changes

Bond energies and enthalpy changes

99

101

Equilibrium

Reversible reactions and equilibrium

117

119

5

States of matter

The liquid state

77

78

9

Rates of reaction

Collision theory

141

142

O

Present?

W

Syllabus content

EV IE

Section number

LY

Cambridge International Pre-University Chemistry

(a) recall the relationship between shell number and number of subshells

Y

B

3

Electrons in atoms

Subshells and atomic orbitals

37

37

S3

S3.5

Subshells and atomic orbitals

A1.2b

(b) recognise an orbital as a mathematical function giving rise to a probability distribution, as illustrated by orbitals such as 1s, 2s, 2px etc. (no knowledge of orbitals from f or higher subshells is required)

Y

B

3

Electrons in atoms

Shapes of the orbitals

38

38

S3

S3.6

Orbitals and the Periodic Table

A1.2c

(c) draw the approximate shapes of the s, p and d orbitals

Y

B

3

Electrons in atoms

Shapes of the orbitals

38

38

S3

S3.7

Atomic orbitals

A1.2d

(d) describe qualitatively the relative energies of the s, p, d and f subshells within a principal shell and describe the concept of shielding in multi-electron atoms

Y

B

3

Electrons in atoms

Filling the shells and orbitals

38

38

S3

S3.7

Atomic orbitals

FO R

A1.2a

Original material © Cambridge University Press 2016

Chapter title

A1.2e

(e) recall that all electrons possess spin and that electrons are spin-paired in orbitals

Y

B

3

Electrons in atoms

A1.2f

(f) apply the aufbau principle

Y

S

A1.2g

(g) state the ground state electronic configurations of the first 36 elements and any common ion using conventional notation and shorthand (e.g. [Ar] 4s2 for Ca) given proton number and charge

Y

B

3

A1.2h

(h) describe the ground state configuration of electrons in a subshell using electron-in-box notation

Y

B

3

A1.2i

(i) explain the variation of atomic radii, ionic radii, first ionisation energies and electronegativities down groups and across the third period in terms of shell number, nuclear charge and shielding; explain the effect of ionisation on atomic radius

Y

B

3

A1.2j

(j) understand and use the relationship between group number and successive ionisation energies; write equations for successive ionisation energies.

Y

B

Supplement

Section number

Start page

End page

Filling the orbitals

40

41

Electrons in atoms

FO R

S3.8

Electronic configurations of atoms – ‘filling’ orbitals with electrons

S3

S3.8

Electronic configurations of atoms – ‘filling’ orbitals with electrons

S3.8

Electronic configurations of atoms – ‘filling’ orbitals with electrons

Filling the orbitals

40

41

S3

S3.8

Electronic configurations of atoms – ‘filling’ orbitals with electrons

Electrons in atoms

Patterns in ionisation energies in the Periodic Table

41

42

S3

S3.9

The trend of first ionisation energies across a period

3

Electrons in atoms

Evidence for electronic structure

34

37

S3

S3.1, S3.2, S3.3, S3.4

Simple electronic structure; Evidence for electronic structure; Three factors that influence the successive ionisation energies of an atom; Interpreting successive ionisation

4

Chemical bonding

Types of chemical bonding; Ionic bonding; Covalent bonding

49

54

Electrons in atoms

• molecular shape and bond angles • intermolecular forces

students should be able to:

C

S3

S3

• σ and π bonds, bond order

Y

Title

40

• molecular bonding and antibonding orbitals

(a) use dot-cross diagrams to describe ionic and covalent bonding, including dative covalent bonding

Section number

38

Content

A1.3a

To add to which chapter of Coursebook?

Electronic configurations

R

A1.3 Chemical forces

Section title

N

Coursebook Chapter number

O

Location (C)oursebook or (S)upplement or (B)oth

A1.3

Present?

W

Syllabus content

EV IE

Section number

LY

Cambridge International Pre-University Chemistry

Original material © Cambridge University Press 2016

Section number

Syllabus content

Location

Coursebook

(C)oursebook or (S)upplement or (B)oth

Chapter number

Supplement

A1.3b

(b) understand that electrons occupy orbitals in molecules as well as atoms; understand that putting electrons into a bonding orbital leads to the nuclei being held together (i.e. bonded) whereas putting electrons into an antibonding orbital weakens bonding by pulling the nuclei apart from one another (as these orbitals lie outside the internuclear region) (N.B. there is no requirement to appreciate the formation of molecular orbitals from the linear combination of atomic orbitals)

Y

S

A1.3c

(c) understand that the strength of the covalent bond can be related to the extent to which the energy of the bonding molecular orbital is lowered relative to the atomic orbitals

Y

S

A1.3d

(d) understand how σ and π bonds result from overlap of atomic orbitals

Y

C

4

A1.3e

(e) appreciate the relative strengths of σ and π bonds

Y

C

4

A1.3f

(f) understand the concept of bond order and its qualitative relationship to bond length and strength

Y

B

4

A1.3g

(g) deduce the shapes and bond angles of simple molecules and ions (typified by examples such as BF3, CH4, NH3, H2O, CO2, C2H4, CH3) and hypervalent species (typified by examples such as PCl5, SF6, IF 7, XeO 4, PCl4) using the VSEPR model

Y

C

A1.3h

(h) deduce changes in geometry and bond angle during chemical reaction, e.g. protonation of ammonia, addition to double bonds

Y

S

A1.3i

(i) describe the concept of electronegativity and its use in predicting bond dipoles

Y

C

4

Chemical bonding

Intermolecular forces

60

63

A1.3j

(j) describe van der Waals forces (all intermolecular forces weaker than hydrogen bonds) in terms of instantaneous, permanent and induced dipoles and the increase of these forces with polarisability

Y

C

4

Chemical bonding

Intermolecular forces

60

63

A1.3k

(k) describe hydrogen bonding as a special case of a dipoledipole interaction and its influence on the structure of ice and the anomalous properties of water; outline its importance in base-pairing in the double helix of DNA (structure of DNA bases is not required)

Y

B

4

Chemical bonding

Hydrogen bonding

64

66

A1.3l

(l) understand the importance of van der Waals forces and hydrogen bonding in determining protein structure (recall of alpha helix and beta sheet structures is not required).

Y

S

A1.4

A1.4 Energy changes

Section number

R

FO R

Section title

Start page

End page

O

N

Chapter title

σ bonds and π bonds

57

Chemical bonding

σ bonds and π bonds

57

57

Chemical bonding

Bond length and bond energy

54

54

Shapes of molecules; More molecular shapes

55

56

W

Chemical bonding

EV IE

Present?

LY

Cambridge International Pre-University Chemistry

4

Chemical bonding

Content

• standard enthalpy changes and the link with temperature • Hess’s law and Born-Haber cycles

• reaction pathway diagrams, including the effect of catalysis

Original material © Cambridge University Press 2016

To add to which chapter of Coursebook?

Section number

Title

S4

4.1

Chemical bonding

S4

S4.1

Chemical bonding

S4

S4.2

Bond order

S4

S4.3

Changes in Geometry

S4

S4.4

Intermolecular forces in DNA

S4

S4.5

Protein structure

57

Syllabus content

Present?

Location

Coursebook

(C)oursebook or (S)upplement or (B)oth

Chapter number

Chapter title

Supplement

Section number

students should be able to: (a) define and use the terms standard enthalpy change of reaction, formation, combustion, hydration, solution, neutralisation, atomisation and vaporisation

Y

C

6

Enthalpy changes

A1.4b

(b) use the relationship q = mc ΔT (Equation 1 in the Data Booklet)

Y

C

6

Enthalpy changes

(c) state Hess’s law and the first law of thermodynamics

Y

C

6

Enthalpy changes

(d) apply Hess’s law to the indirect determination of enthalpy changes, including the use of average gas- phase bond energies

Y

C

6

Enthalpy changes

A1.4e

(e) calculate lattice energies from Born-Haber cycles

A1.4f

(f) understand the effect of a catalyst in providing an alternative reaction route with a lower activation energy

A1.4g

(g) interpret and construct reaction profile energy diagrams, including the effect of a catalyst on the activation energy, and deduce the different energy changes involved

Y

A1.4h

(h) interpret the effect of a catalyst in terms of the Boltzmann distribution.

Y

B1.5

B1.5 Free energy and entropy

EV IE

W

A1.4c A1.4d

Standard enthalpy changes

O

A1.4a

Start page

End page

Measuring enthalpy changes

92

93

94

96

Hess’s law

97

97

Hess’s law; Enthalpy change of reaction from enthalpy changes of formation; Enthalpy change of formation from enthalpy changes of combustion; Calculating the enthalpy change of hydration of an anhydrous salt; Bond energies and enthalpy changes; Calculating enthalpy changes using bond energies

97

101

Y

B

19

Lattice energy

[This syllabus point links with all of coursebook chapter 19]

258

269

Y

C

9

Rates of reaction

Catalysis

144

145

C

9

Rates of reaction

Reaction kinetics

141

142

C

9

Rates of reaction

The effect of temperature on rate of reaction

143

144

R

FO R

Section title

To add to which chapter of Coursebook?

Section number

Title

S19

S19.1

Limitations of lattice energy calculations

N

Section number

LY

Cambridge International Pre-University Chemistry

Content

• entropy and standard entropy change • the second law of thermodynamics Part B

• the Gibbs equation and the equilibrium constant

Original material © Cambridge University Press 2016

Syllabus content

Present?

Location

Coursebook

(C)oursebook or (S)upplement or (B)oth

Chapter number

Chapter title

Supplement

Section number

Y

C

23

Entropy and Gibbs free energy

B1.5b

(b) understand that the total entropy change is the entropy change of the system plus the entropy change of the surroundings

Y

C

23

Entropy and Gibbs free energy

B1.5c

(c) calculate Δ S given the standard entropies of the reactants and products

Y

C

23

Entropy and Gibbs free energy

B1.5d

(d) define the entropy change of the surroundings as –ΔrH/ T where T is the thermodynamic temperature of the surroundings in Kelvin and ΔrH is the enthalpy change of the reaction

Y

C

23

B1.5e

(e) recall and apply the second law of thermodynamics – that the total entropy (i.e. system plus surroundings) increases in every chemical reaction until equilibrium is reached – and explain how it can account for endothermic phenomena

Y

C

23

B1.5f

(f) use the Gibbs equation (Equation 2 in the Data Booklet)

B1.5g

(g) use the equation that relates the standard Gibbs energy change to the equilibrium constant (Equation 3 in the Data Booklet).

B1.6

B1.6 Equilibrium • equilibrium constants and Le Chatelier’s principle

• pH and buffers • oxidation number; quantitative electrolysis • standard electrode and cell potentials • electron flow in cells; Gibbs energy change

C

Y

S

23

354

Calculating entropy changes

354

357

Calculating entropy changes

354

357

Entropy and Gibbs free energy

Entropy and temperature

357

357

Entropy and Gibbs free energy

Entropy, enthalpy changes and free energy

357

358

Gibbs free energy; Gibbs free energy calculations

358

362

Entropy and Gibbs free energy

To add to which chapter of Coursebook?

Section number

Title

S23

S23.1

The standard Gibbs energy change and the equilibrium constant

S8

S8.1 S8.2

The Lewis theory of acids and bases Mono-, di- and tri-protic acids

R

• Brønsted-Lowry and Lewis theories of acids and bases

350

EV IE

Content

Y

End page

Introducing entropy; Chance and spontaneous change

W

(a) understand entropy in terms of: (i) the random dispersal of molecules in space, and (ii) the random dispersal of quanta of energy among molecules

Start page

O

students should be able to: B1.5a

Section title

N

Section number

LY

Cambridge International Pre-University Chemistry

FO R

students should be able to: B1.6a

(a) state Le Chatelier’s principle and apply it to deduce qualitatively the effects of changes in temperature, concentration or pressure on a system at equilibrium

Y

C

8

Equilibrium

Changing the position of equilibrium

119

123

B1.6b

(b) describe weak acids and alkalis in terms of equilibria; describe the Brønsted-Lowry theory of acids, including conjugate pairs; describe the Lewis theory of acids and bases; understand the difference between mono-, di- and triprotic (tribasic) acids

Y

B

8

Equilibrium

Acid–base equilibria

130

133

Original material © Cambridge University Press 2016

Syllabus content

Present?

Location

Coursebook

(C)oursebook or (S)upplement or (B)oth

Chapter number

Chapter title

B1.6c

(c) understand that equilibrium constant expressions take the form outlined in Equation 4 of the Data Booklet , and apply this to K c

Y

B

8

Equilibrium

B1.6d

(d) understand that in the case of K p only gas phase constituents are included in the equilibrium constant expression, that in the case of K sp solids are omitted, and that in the case of K a and K w the solvent is omitted

Y

B

8

Equilibrium

B1.6d

(d) understand that in the case of K p only gas phase constituents are included in the equilibrium constant expression, that in the case of K sp solids are omitted, and that in the case of K a and K w the solvent is omitted

Y

C

21

Further aspects of equilibria

B1.6d

(d) understand that in the case of K p only gas phase constituents are included in the equilibrium constant expression, that in the case of K sp solids are omitted, and that in the case of K a and K w the solvent is omitted

Y

C

21

B1.6e

(e) calculate quantities based on equilibrium constant expressions

Y

C

8

B1.6f

(f) define pH and determine pH for strong and weak acids (solving of quadratic equations is not required); understand how pK a is a measure of acid strength

Y

C

B1.6g

(g) describe and explain the change in pH during strong/weak acid-base titrations and explain the choice of indicator

Y

B1.6h

(h) understand and explain how simple buffer solutions work; calculate the pH of buffer solutions

B1.6i

(i) explain the electrolysis of a molten compound using inert electrodes; predict the products of the electrolysis of an aqueous electrolyte, given relevant electrode potentials; construct half-equations for the reactions occurring at each electrode during electrolysis

B1.6i

(i) explain the electrolysis of a molten compound using inert electrodes; predict the products of the electrolysis of an aqueous electrolyte, given relevant electrode potentials; construct half-equations for the reactions occurring at each electrode during electrolysis

Y

B1.6j

(j) calculate quantities of substances involved during electrolysis using the Faraday constant (given in the Data Booklet)

B1.6k B1.6l

Supplement

Section number

Section title

Start page

End page

N

Section number

LY

Cambridge International Pre-University Chemistry

To add to which chapter of Coursebook?

Section number

Title

123

126

S8

S8.3

What are the units of K c?

Equilibria in gas reactions: the equilibrium constant, K p

127

128

S8

S8.4

What are the units of K p?

Equilibrium and solubility

316

319

Further aspects of equilibria

The ionic product of water, K w

304

304

Equilibrium

Some examples of equilibrium calculations

124

126

21

Further aspects of equilibria

pH calculations; Weak acids – using the acid dissociation constant, K a

305

309

C

21

Further aspects of equilibria

Indicators and acid– base titrations

309

313

Y

C

21

Further aspects of equilibria

Buffer solutions

313

315

Y

C

20

Electrochemistry

Electrolysis

275

275

C

20

Electrochemistry

Electrode potentials

278

281

Y

C

20

Electrochemistry

Quantitative electrolysis

276

287

(k) understand the concept of redox with reference to oxidation number

Y

C

7

Redox reactions

Redox and oxidation number

110

111

(l) understand the concept of standard electrode potential and describe the standard hydrogen electrode

Y

C

20

Electrochemistry

The standard hydrogen electrode

279

279

FO R

R

EV IE

W

O

Equilibrium expressions and the equilibrium constant, K c

Original material © Cambridge University Press 2016

Coursebook Chapter number

Chapter title

B1.6m

(m) describe methods used to measure standard electrode potentials of: (i) metals or non-metals in contact with their ions in aqueous solution, and (ii) ions of the same element in different oxidation states

Y

C

20

Electrochemistry

B1.6n

(n) understand and be able to use standard notation to construct and interpret conventional cell diagrams, e.g. Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s); Pt | H (g) | 2H+(aq) || MnO –(aq) + 8H+(aq), Mn2+(aq) + 4H2O(l) | Pt; be able to relate cell diagrams to pictorial diagrams of cells

Y

S

B1.6o

(o) calculate standard cell potential from standard electrode potentials

Y

C

B1.6p

(p) understand the link between the standard cell potential and the standard Gibbs energy change and use Equation 5 in the Data Booklet

Y

S

B1.6q

(q) deduce the direction of electron flow in an electrochemical cell and therefore the feasible direction of reaction using cell potentials

Y

C

B1.6r

(r) understand how the cell potential changes from its standard value during the course of a reaction in an electrochemical cell and approaches zero as equilibrium is approached; interpret the variation of the cell potential as the concentrations of the solutions change using Le Chatelier’s principle (knowledge of the Nernst equation is not required)

Y

C

20

B1.6s

(s) describe the hydrogen/oxygen fuel cell.

Y

C

B1.7 Gases and kinetics Content

• homogeneous and heterogeneous catalysis

Electrochemistry

Section title

Start page

End page

Measuring standard electrode potentials

Using E values

282

284

Title

S20

S20.1

Standard notation and cell diagrams

S20

S20.2

Standard cell potentials and Gibbs energy changes

S5

S5.1

Boyle’s law and Charles’s law

285

285

290

Electrochemistry

How does the value of E vary with ion concentration?

290

291

20

Electrochemistry

Hydrogen–oxygen fuel cells

294

295

Electrochemistry

Section number

284

Using E values

20

To add to which chapter of Coursebook?

R

• kinetic theory and ideal gases • rate of reaction, Arrhenius equation

20

Supplement

Section number

N

Location (C)oursebook or (S)upplement or (B)oth

O

Present?

W

Syllabus content

EV IE

Section number

LY

Cambridge International Pre-University Chemistry

• formulating rate equations from data; multi-step reaction mechanisms

FO R

• rate-concentration and concentration-time dependence students should be able to: B1.7a

(a) state the assumptions of the kinetic theory applied to an ideal gas, and understand their limitations

Y

C

5

States of matter

The gaseous state

73

74

B1.7a

(a) state the assumptions of the kinetic theory applied to an ideal gas, and understand their limitations

Y

C

5

States of matter

Limitations of the ideal gas laws

75

75

B1.7b

(b) state Boyle’s law and Charles’s law and use the ideal gas equation, pV = nRT (Equation 6 in the Data Booklet)

Y

B

5

States of matter

Ideal gases

74

75

B1.7b

(b) state Boyle’s law and Charles’s law and use the ideal gas equation, pV = nRT (Equation 6 in the Data Booklet)

Y

C

5

States of matter

The general gas equation

75

76

Original material © Cambridge University Press 2016

Chapter title

(c) explain the dependence of the rate of reaction on the activation energy of the reaction and temperature in terms of the Boltzmann distribution of molecular energies

Y

C

9

Rates of reaction

B1.7d

(d) rationalise the rate constant as the product of collision frequency (per unit concentration), the steric factor and a Boltzmann term; use the Arrhenius equation (Equation 7 in the Data Booklet)

Y

S

B1.7e

(e) describe the homogeneous and heterogeneous modes of catalytic activity, and that it is the shape of the active site in enzymes that is central to their function and makes them highly specific

Y

C

9

B1.7f

(f) understand how rate equations are based on experimental data; deduce the orders and overall order of reaction from a rate equation; deduce the molecularity of a given step; deduce the units of rate of reaction, and of rate constants for different orders of reaction

Y

C

22

B1.7g

(g) interpret simple multi-step reactions in terms of a series of elementary steps with a rate-determining step; predict orders and overall order of reaction from a given mechanism

Y

C

22

B1.7h

(h) suggest a mechanism that is consistent with a given rate equation

Y

C

B1.7i

(i) deduce rate equations using initial rates data; interpret rateconcentration graphs

Y

B1.7j

(j) understand that only first-order reactions have a constant half life; use the equation for the concentration-time dependence for first-order reactions (Equation 8 in the Data Booklet) to calculate the first-order rate constant and half lives from provided data.

Y

Content

Rates of reaction

Section title

The effect of temperature on rate of reaction

End page

143

145

Rate equations; Which order of reaction?

Calculations involving the rate constant, k

330

335

Reaction kinetics

Kinetics and reaction mechanisms

338

340

22

Reaction kinetics

Verifying possible reaction mechanisms

338

340

C

22

Reaction kinetics

Deducing order of reaction from raw data

335

338

B

22

Reaction kinetics

Half-life and reaction rates

333

335

Reaction kinetics

• the limitations of scientific models

Original material © Cambridge University Press 2016

To add to which chapter of Coursebook?

Section number

Title

S22

S22.1

The Arrhenius Equation

S22

S22.2

Calculating k for first order reactions

144

144

• the relationship between evidence, models and theories

FO R

Start page

Catalysis; Enzymes

W

B1.7c

Supplement

Section number

N

Coursebook Chapter number

O

Location (C)oursebook or (S)upplement or (B)oth

B1.8 Chemical models and evidence

Present?

EV IE

Syllabus content

R

Section number

LY

Cambridge International Pre-University Chemistry

Syllabus content

Present?

Location

Coursebook

(C)oursebook or (S)upplement or (B)oth

Chapter number

Chapter title

Supplement

C

14

Introduction to organic chemistry

16

Halogenoalkanes

22

Reaction kinetics

Section number

students should be able to: Y

22

Start page

End page

Organic reactions – mechanisms

196

198

Mechanism of nucleophilic substitution in halogenoalkanes

220

222

Deducing order of reaction from raw data

335

337

Reaction kinetics

Verifying possible reaction mechanisms

338

340

States of matter

The gaseous state

73

77

States of matter

The gaseous state

73

77

O

(a) appreciate that chemistry is an evidence-based subject. The experimental basis, in particular, of rate equations and organic mechanisms provide useful examples (see Section 3, Organic chemistry)

W

B1.8a

Section title

To add to which chapter of Coursebook?

Section number

Title

N

Section number

LY

Cambridge International Pre-University Chemistry

(b) understand the distinction between theories, models, facts and definitions

Y

C

5

B1.8c

(c) recognise the ideal gas model and equation as an example of a model based on a series of assumptions and appreciate that it is useful up to a point but breaks down under certain conditions

Y

C

5

B1.8d

(d) recognise the limitations (increasing with covalent character) of the ionic model of crystal lattices to predict lattice energies found from Born-Haber cycles.

Y

B

19

Lattice energy

Born–Haber cycles

259

262

S19

S19.1

Limitations of lattice energy calculations

S10

S10.1

Three types of elements

2

EV IE

B1.8b

Inorganic chemistry

This section is intended to give students understanding of the three main types of bonding within the framework of electronegativity and in the continuum of the van Arkel diagram; to enable students to rationalise a body of descriptive chemistry in terms of group and periodic trends; knowledge and understanding of the structure, isomerism and redox properties of transition metal complexes; and appreciation of threedimensional relationships in crystal structures.

A2.1 Periodic Table Content Part A

R

A2.1

• division of the Periodic Table by element and structure type

FO R

• rationalisation and prediction of bonding type using the van Arkel triangle • periodic trends in physical and chemical properties students should be able to: A2.1a

(a) appreciate that the elements can be divided into metals, metalloids and non-metals, and the locations of these regions on the Periodic Table

Y

B

10

Periodicity

3 Periodic patterns of melting points and electrical conductivity

151

153

A2.1b

(b) appreciate that elemental structures can be divided into simple and giant

Y

C

5

States of matter

States of matter; Giant molecular structures

73

81

Original material © Cambridge University Press 2016

Coursebook Chapter number

Chapter title

(c) appreciate that some elements exist as different allotropes; recall the allotropes of oxygen; describe the structure and properties of the allotropes of carbon, including buckminsterfullerene and graphene

Y

B

5

States of matter

A2.1c

(c) appreciate that some elements exist as different allotropes; recall the allotropes of oxygen; describe the structure and properties of the allotropes of carbon, including buckminsterfullerene and graphene

Y

C

5

States of matter

A2.1d

(d) understand the variation of melting and boiling points and electrical conductivity of elements across Period 3 in terms of structure and bonding

Y

B

10

Periodicity

A2.1e

(e) predict the bonding type, and hence properties, in elements and compounds using electronegativity values and the van Arkel triangle; plot points on a template van Arkel triangle

Y

B

10

A2.1e

(e) predict the bonding type, and hence properties, in elements and compounds using electronegativity values and the van Arkel triangle; plot points on a template van Arkel triangle

Y

B

10

A2.1f

(f) deduce unfamiliar half-equations and construct redox equations from the relevant half-equations

Y

C

7

A2.1g

(g) understand and apply the concept of oxidation number (including use in nomenclature) and appreciate that the maximum oxidation states of elements are found in compounds with oxygen and fluorine

Y

B

A2.1h

(h) describe the reactivity of Period 3 elements with oxygen and water

Y

C

A2.1i

(i) describe the reactivity of Period 3 oxides with water, and the acid-base character of the oxides as the periodic table is crossed

Y

A2.1j

(j) describe the reactivity of Period 3 chlorides with water, including reactions of SiCl4 and PCl5 with water in terms of replacement of –Cl with –OH and subsequent dehydration; analogy with dehydration of unstable hydroxyl group in organic chemistry

A2.1k

(k) appreciate that there is a thermodynamic preference of bonds to oxygen compared to chlorine, as illustrated by the natural occurrence of mineral oxides.

Start page

End page

To add to which chapter of Coursebook?

Section number

Title

S5

S5.2

The allotropes of oxygen

Giant molecular structures

80

81

Carbon nanoparticles

82

83

3 Periodic patterns of melting points and electrical conductivity

151

153

S10

S10.2

Periodic patterns of structure in metals

Periodicity

Oxides of Period 3 elements

156

158

S10

S10.3

The van ArkelKetelaar triangle

Periodicity

Chlorides of Period 3 elements

158

160

S10

S10.4

More about the effect of water on chlorides of Period 3 elements

Redox reactions

What is a redox reaction?; Redox and electron transfer

107

109

7

Redox reactions

Oxidation numbers

109

113

S7

S7.1

Oxidation numbers

10

Periodicity

React ions of Period 3 elements with oxygen; React ions of sodium and mag nesium with water

154

155

C

10

Periodicity

Oxides of Period 3 elements

156

158

Y

C

10

Periodicity

Chlorides of Period 3 elements

158

160

Y

S

S10

S10.4

More about the effect of water on chlorides of Period 3 elements

R

FO R

Section title

O

A2.1c

Supplement

Section number

N

Location (C)oursebook or (S)upplement or (B)oth

A2.2

Present?

W

Syllabus content

EV IE

Section number

LY

Cambridge International Pre-University Chemistry

A2.2 Main group chemistry (except Group 14) Content

• thermal stability of carbonates

• properties of nitrogen, phosphorus and ammonia

Original material © Cambridge University Press 2016

Syllabus content

Present?

Location

Coursebook

(C)oursebook or (S)upplement or (B)oth

Chapter number

Chapter title

Supplement

Section number

• chemistry of oxygen and sulfur • halogen chemistry and the exceptional behaviour of fluorine

Y

C

11

Group 2

A2.2a

(a) describe the thermal decomposition reactions of carbonates, and explain the trends in thermal stability in terms of the charge density of the cation and the polarisability of the anion

Y

C

19

Lattice energy

A2.2b

(b) recall and compare the structures of nitrogen and white phosphorus

Y

B

13

A2.2c

(c) explain why nitrogen is unreactive

Y

S

A2.2d

(d) describe the acid-base behaviour of ammonia and the ammonium ion

A2.2e

(e) describe the redox properties of hydrogen peroxide

A2.2f

W

(a) describe the thermal decomposition reactions of carbonates, and explain the trends in thermal stability in terms of the charge density of the cation and the polarisability of the anion

End page

To add to which chapter of Coursebook?

Section number

Title

168

168

The thermal stability of Group 2 carbonates and nitrates

264

264

EV IE

Nitrogen and sulfur

Nitrogen gas

181

182

S13

S13.1

White phosphorus

S13

S13.2

Why is nitrogen unreactive?

S13

S13.5

Hydrogen peroxide: a redox chameleon

C

Y

S

(f) describe the formation of sulfur oxides in the atmosphere from carbon fuels containing sulfur, and the removal of SO2 from power station emissions

Y

B

13

Nitrogen and sulfur

Sulfur and its oxides

185

185

S13

S13.3

Reducing sulfur emissions

A2.2g

(g) describe sulfuric acid as a strong involatile acid, oxidising agent, dehydrating agent and catalyst; understand the importance of sulfuric acid as an industrial chemical

Y

B

13

Nitrogen and sulfur

Sulfuric acid

185

185

S13

S13.4

Reactions of sulfuric acid

A2.2h

(h) recall the trends in volatility and colour, and understand the trend in oxidising power of the halogens, including fluorine

Y

C

12

Group 17

Physical properties of Group 17 elements

172

172

A2.2i

(i) explain the trend in bond energy of the halogens, with the reason for F2 being exceptional

Y

B

12

Group 17

Reactions of Group 17 elements

173

173

S12

S12.1

Trend in bond enthalpies

A2.2j

(j) explain the trend in acidity of the hydrogen halides in terms of the bond energy, including HF being a weak acid

Y

B

12

Group 17

Reactions with hydrogen

174

175

S12

S12.2

Acidity of hydrogen halides

A2.2k

(k) recall the reaction of halides with concentrated (conc.)

H2SO 4 and explain the trend towards oxidation

Y

B

12

Group 17

Reactions of halide ions with concentrated sulfuric acid

176

177

S12

S12.3

Oxidation of hydrogen halides

A2.2l

(l) recall the reaction of iodine with sodium thiosulfate and its use in analysis

Y

B

12

Group 17

Uses of the halogens and their compounds

178

178

S12

S12.4

Reaction between iodine and the thiosulfate ion

A2.2m

(m) recall the reactions of the halogens with cold NaOH(aq) in terms of disproportionation

Y

C

12

Group 17

Disproportionation

177

177

A2.2n

(n) state reasons for the anomalous reactivity of fluorine (see above), and understand why elements can form high-oxidation state fluorides, such as UF6, SF6.

Y

B

12

Group 17

Reactions of Group 17 elements

173

173

S12

S12.1

Trend in bond enthalpies

R

Y

FO R

13

Start page

Thermal decomposition of Group 2 carbonates and nitrates

O

students should be able to: A2.2a

Section title

N

Section number

LY

Cambridge International Pre-University Chemistry

Nitrogen and sulfur

Ammonia and ammonium compounds

Original material © Cambridge University Press 2016

182

183

Syllabus content

Present?

Location

Coursebook

(C)oursebook or (S)upplement or (B)oth

Chapter number

Supplement

Chapter title

Section number

Part B B2.3

Section title

Start page

End page

To add to which chapter of Coursebook?

Section number

Title

S31

S31.1

Group 14

S31

S31.1

Group 14

S31

S31.1

Group 14

S31

S31.1

Group 14

N

Section number

LY

Cambridge International Pre-University Chemistry

B2.3 From non-metals to metals: Group 14 Content

O

• trends in the properties of the elements • properties of the oxides

Y

S

B2.3b

(b) recall that lead chemistry is dominated by the +2 oxidation state

Y

S

B2.3c

(c) recall that the oxides are essentially covalent at the top of the group and largely ionic at the bottom

Y

S

B2.3d

(d) explain the bonding in the divalent and tetravalent oxides, making reference to the van Arkel diagram and polarisation.

Y

S

B2.4

B2.4 Transition elements

Y

B

Content • physical and atomic properties • geometry and isomerism of complexes • colour and ligand substitution • redox chemistry of complexes students should be able to:

EV IE

(a) describe the transition from non-metal through metalloid to metal, and its manifestation in the electrical conductivities of the elements

W

students should be able to: B2.3a

(a) explain what is meant by a transition element

24

Transition elements

What is a transition element?

367

369

S24

S24.1

More about physical properties of the transition elements

B2.4b

(b) state a qualitative comparison between the properties of a transition element and an s-block element in terms of atomic radius, melting and boiling points and first ionisation energy

Y

B

24

Transition elements

Physical properties of the transition elements

369

369

S24

S24.1

More about physical properties of the transition elements

B2.4c

(c) state that there is little variation in the above properties across the series compared to the rest of the period

Y

B

24

Transition elements

Physical properties of the transition elements

369

369

S24

S24.1

More about physical properties of the transition elements

B2.4d

(d) recall the characteristic properties of transition elements, including variable oxidation states, formation of stable complexes and catalysis, using common examples

Y

B

24

Transition elements

What is a transition element?

367

369

S24

S24.1

More about physical properties of the transition elements

B2.4d

(d) recall the characteristic properties of transition elements, including variable oxidation states, formation of stable complexes and catalysis, using common examples

Y

B

24

Transition elements

Ligands and complex formation

371

373

S24

S24.2

More about ligands and complex formation

FO R

R

B2.4a

Original material © Cambridge University Press 2016

Coursebook Chapter number

Chapter title

Y

B

24

Transition elements

B2.4f

(f) explain the isomerism in complexes: geometric and optical

Y

B

24

Transition elements

B2.4g

(g) state that the d orbitals point either along or between the cartesian axes, which explains a splitting of energy levels in transition metal complexes

Y

B

24

Transition elements

B2.4h

(h) explain qualitatively the origin of colour in transition metal complexes

Y

B

24

B2.4i

(i) understand that different ligands give different coloured complexes and hence colour changes are often observed during ligand exchange

Y

B

24

B2.4j

(j) describe the following redox chemistry: Fe3+/Fe2+; MnO –/ Mn2+; Cr2O72–/Cr3+; and Cu2+/Cu+

Y

B

B2.4k

(k) outline the essential biological role of the following iron complexes: haemoglobin, myoglobin and ferritin

Y

S

B2.4l

(l) recall that a ligand may affect the relative stability of oxidation states of a transition metal complex, e.g. the +2 and +3 oxidation states of cobalt: [Co(H2O)6]2+/[Co(H2O)6]3+ versus [Co(NH3)6]2+/[Co(NH3)6]3+.

Y

B

B2.5

B2.5 Crystal structures • close packing in metals • unit cell properties • occupying holes in unit cells

24

24

Section title

Ligands and complex formation

Start page

End page

To add to which chapter of Coursebook?

Section number

Title

371

373

S24

S24.2

More about ligands and complex formation

Stereoisomerism in transition metal complexes

373

373

S24

S24.3

More about isomerism

The colour of complexes

376

378

Transition elements

The colour of complexes

376

378

S24

S24.4

More about substitution of ligands

Transition elements

The colour of complexes

376

378

S24

S24.5

More about the colour of complexes

Redox reactions

369

371

S24

S24.6

Redox reactions

S24

S24.7

Transition metal complexes in biological systems

S24

S24.8

Effect of ligands on complex stability

O

(e) describe the geometry and bond angles around the following transition metal ions in complexes: octahedral hexaaqua ions of first row transition metals, tetrahedral complexes of these ions with larger ligands (e.g. [CoCl4]), and square planar complexes with Group 10 metals, especially platinum (e.g. Pt(NH3)2Cl2)

W

B2.4e

Supplement

Section number

N

Location (C)oursebook or (S)upplement or (B)oth

Content

Present?

EV IE

Syllabus content

Transition elements

Transition elements

Stability constants

375

376

R

Section number

LY

Cambridge International Pre-University Chemistry

students should be able to: (a) describe and recognise the cubic close-packed (CCP) and hexagonal close-packed (HCP) structures in metals, including the ABC and AB representations of the close-packed structures

Y

S

S32

S32.2

Close-packing structures

B2.5b

(b) define the unit-cell representation as the simplest repeating unit of the lattice which displays the full symmetry of the crystal

Y

S

S32

S32.3

The unit cell

B2.5c

(c) explain the relationship between neighbouring atoms/ions in a lattice in terms of geometry (limited to tetrahedral and octahedral) and coordination number

Y

S

S32

S32.3

The unit cell

B2.5d

(d) appreciate that there are octahedral and tetrahedral holes in close-packed structures and recall the ratio of holes to atoms in each case

Y

S

S32

S32.4

From metals to compounds in close-packed structures

FO R

B2.5a

Original material © Cambridge University Press 2016

B2.5e

(e) understand the derivation of lattice structures by occupying holes in a close-packed lattice of ions with counter-ions: NaCl by filling all the octahedral holes in the CCP lattice, CaF2 from filling all the tetrahedral holes

Y

S

B2.5f

(f) explain that lattice energies may be calculated from the crystal structure using electrostatics and the ionic model, and the deficiencies of this method (recall of the Born-Landé equation is not required).

Y

S

3

Present?

Location

Coursebook

(C)oursebook or (S)upplement or (B)oth

Chapter number

Supplement

Chapter title

Section number

Section title

Start page

End page

N

Syllabus content

O

Section number

LY

Cambridge International Pre-University Chemistry

Organic chemistry

W

This section is intended to equip students with the tools to understand organic reactions through the framework of the functional group level, to suggest reagents and conditions for the transformations they have studied when encountered in unfamiliar reaction schemes, to consider the three-dimensional nature of organic reactions, and to understand mechanisms of organic reactions and the acid-base properties of organic molecules.

EV IE

Greater emphasis is placed on understanding what kind of reagent is required for a chemical transformation,

A3.1 Preliminaries Content • formulae, structures and geometry • structural, geometric and optical isomerism • nomenclature • types of reaction Part A students should be able to:

(a) interpret and use the terms: • molecular formula, as the actual number of atoms of each element in a molecule, e.g. C 3H8O for propan-1-ol, not C 3H7OH; • general formula, as the simplest algebraic formula of a member of a homologous series, e.g. C n H2n + 2 for an alkane; • structural formula, as the minimal detail that shows the arrangement of atoms in a molecule, e.g. CH3CH2CH2OH for propan-1-ol, not C 3H7OH; • displayed formula, as the relative positioning of atoms and the bonds between them, e.g. for ethanol: H H

FO R

A3.1a

R

rather than the learning of reagents and conditions for all reactions. Those reactions where specific reagents are required are detailed in the learning outcomes; in other cases only the kind of reagent (e.g. oxidising agent) and an example of each kind of reagent will be expected.

H

C

C

H

H

O

Y

C

14

Introduction to organic chemistry

Representing organic molecules

H

Original material © Cambridge University Press 2016

189

191

To add to which chapter of Coursebook?

Section number

Title

S32

S32.4

From metals to compounds in close-packed structures

S32

S32.6

A theoretical treatment of ionic bonding (the Born–Landé equation)

Syllabus content

Present?

Location

Coursebook

(C)oursebook or (S)upplement or (B)oth

Chapter number

H3C H

H

H

convention for representing the aromatic

ring is preferred.

EV IE

HO

The

Section title

Start page

End page

W

H

CH3

Section number

O

â&#x20AC;˘ skeletal (or partial skeletal) formula, as the simplified organic formula, shown by removing hydrogen atoms from alkyl chains, leaving just a carbon skeleton and associated functional groups, OH e.g. for butan-2-ol (skeletal): and for cholesterol (partial skeletal):

CH3

Supplement

Chapter title

N

Section number

LY

Cambridge International Pre-University Chemistry

(b) appreciate the tetravalent nature of carbon and the shapes of the ethane, ethene and benzene molecules; use this to write structural formulae and predict the shapes of, and bond angles in, other related organic molecules

Y

C

14

Introduction to organic chemistry

Bonding in organic molecules

193

194

A3.1b

(b) appreciate the tetravalent nature of carbon and the shapes of the ethane, ethene and benzene molecules; use this to write structural formulae and predict the shapes of, and bond angles in, other related organic molecules

Y

C

25

Benzene and its compounds

The benzene ring

382

383

A3.1c

(c) draw carbon backbones (i.e. skeletal formulae), showing realistic geometry (zig zag)

Y

C

14

Introduction to organic chemistry

Representing organic molecules

189

191

A3.1d

(d) draw structures and derive molecular formulae from structures (molecular formulae being of the form C aHb X cYd, where heteroatoms are listed in atomic number order)

Y

C

14

Introduction to organic chemistry

Representing organic molecules

189

191

A3.1e

(e) understand structural (including functional group) and geometric (cis-trans) isomerism

Y

C

14

Introduction to organic chemistry

Structural isomerism

194

195

A3.1e

(e) understand structural (including functional group) and geometric (cis-trans) isomerism

Y

C

14

Introduction to organic chemistry

1 Cisâ&#x20AC;&#x201C;trans isomerism

195

195

A3.1f

(f) understand optical isomerism in terms of asymmetric carbons (chiral carbons)

Y

C

14

Introduction to organic chemistry

2 Optical isomerism

195

196

FO R

R

A3.1b

Original material Š Cambridge University Press 2016

To add to which chapter of Coursebook?

Section number

Title

Syllabus content

Present?

A3.1g

(g) use hashed and wedged bonds to represent 3D structures, e.g.

CH3

Coursebook Chapter number

Chapter title

Y

C

14

Introduction to organic chemistry

Introduction to organic chemistry

CH3

Supplement

Section number

Representing organic molecules

C H

H

OH CO2H

CO2H

(h) understand and use nomenclature of molecules with functional groups mentioned in the syllabus, up to six carbon atoms (six plus six for esters)

Y

C

14

A3.1i

(i) use and interpret the terms: oxidation, reduction, hydrolysis, condensation, isomerisation, substitution, addition, elimination, nucleophile/nucleophilic, electrophile/ electrophilic, free-radical, homolytic/heterolytic fission, as appropriate to Part A or Part B; recognise such reactions, as appropriate to Part A or Part B

Y

B

14

A3.1j

(j) understand and use curly arrows to represent movement of electron pairs in reaction mechanisms.

A3.2

A3.2 Functional group level Content • the inactivity of C–H and C–C bonds • the diversity of heteroatom chemistry • the concept of functional group level • moving between functional group levels

students should be able to:

Y

C

R

• unstable groups

Introduction to organic chemistry

EV IE

A3.1h

W

mirror plane

Section title

Start page

End page

189

191

Functional groups; Naming organic compounds

191

193

Organic reactions – mechanisms; Types of organic reaction

196

198

O

C HO

Location (C)oursebook or (S)upplement or (B)oth

14

Introduction to organic chemistry

Organic reactions – mechanisms

196

197

A3.2a

(a) explain the relative inactivity of C–H and C–C single bonds in terms of the absence of lone pairs and dipole moments

Y

C

15

Hydrocarbons

Reactions of alkanes

204

204

A3.2b

(b) explain in terms of dipole moments and lone pairs, why heteroatoms (O, N, X, etc.) lead to more diverse chemistry (dipole moments, high-energy lone pairs)

Y

B

15

Hydrocarbons

Substitution reactions of alkanes

206

207

A3.2c

(c) recognise functional groups (limited to the examples in the syllabus, as appropriate to Part A or Part B) present in a given structure, and name simple molecules containing these groups

Y

C

14

Introduction to organic chemistry

Functional groups; Naming organic compounds

191

193

A3.2d

(d) use the concept of functional group level of a carbon atom from counting the number of bonds to electronegative atoms on the carbon (no such bond is called the Hydrocarbon level; one bond is called the Alcohol level; two bonds is the Carbonyl level; three bonds is the Carboxylic Acid level; and four bonds is the Carbon Dioxide level; appreciate that functional group level applies to a single carbon atom, not the whole molecule

Y

S

FO R

To add to which chapter of Coursebook?

Section number

Title

S14

S14.1

Condensation and isomerisation reactions

S15

S15.1

Diversity of heteroatom chemistry

S15

S15.2

The concept of functional group level

N

Section number

LY

Cambridge International Pre-University Chemistry

Original material © Cambridge University Press 2016

A3.2e

(e) state the functional group level of a carbon atom in an unfamiliar example, and deduce changes in functional group level in a reaction scheme

Y

S

A3.2f

(f) understand that reaction within a level simply swaps one heteroatom for another, e.g. hydrolysis; that moving a carbon atom up a level requires an oxidizing agent, that moving a carbon atom down a level requires a reducing agent or carbanion equivalent; that hydrolysis ultimately yields the functional group after which the level is named (note that not all heteroatoms hydrolyse under typical conditions, e.g. –NH2, –SH, F)

Y

S

A3.2g

(g) recognise common examples within the Alcohol level: alcohols, amines, alkyl halides; the Carbonyl level: aldehydes, ketones; the Carboxylic Acid level: carboxylic acids, esters, amides, nitriles; the Carbon Dioxide level: carbon dioxide

Y

S

A3.2h

(h) describe what happens when there are unstable groups, e.g. two hydroxyls on the same carbon atom (geminal diol or hydrate) forming aldehyde/ketone; three hydroxyls on one carbon atom forming a carboxylic acid; four hydroxyl groups on one carbon atom forming carbonic acid and carbon dioxide.

Location

Coursebook

(C)oursebook or (S)upplement or (B)oth

Chapter number

Supplement

Chapter title

Section number

Section title

Start page

End page

W EV IE

A3.3

Present?

N

Syllabus content

To add to which chapter of Coursebook?

Section number

Title

S15

S15.3

Moving between functional groups

S15

S15.3

Moving between functional groups

S15 S17

S15.2 S17.1 S17.5

The concept of functional group level The alcohol functional group level The carboxylic acid functional group level

S15

S15.4

Unstable groups

S17

S17.1

The alcohol function group level

S17

S17.2

Hydrolysis of fluoroalkanes

O

Section number

LY

Cambridge International Pre-University Chemistry

Y

S

A3.3 Lower functional group level reactions – Alcohol level Content • moving within the level: – synthesis of alcohols and amines from halogenoalkanes – synthesis of halogenoalkanes from alcohols • moving down a level: • moving up a level: – oxidation to aldehydes and ketones students should be able to:

R

– substitution of halogenoalkanes with cyanide

(a) appreciate that the hydrolysis of any member within the group is an example of moving within the functional group level and leads to an alcohol

Y

S

A3.3b

(b) explain how alcohols and amines may be synthesised from halogenoalkanes (using aqueous sodium hydroxide and ethanolic ammonia, respectively), and halogenoalkanes may be synthesised from alcohols using phosphorus halides, and how these are examples of moving within the level

Y

B

16

Halogenoalkanes

1 Substitution reactions with aqueous alkali, OH –(aq)

218

219

A3.3b

(b) explain how alcohols and amines may be synthesised from halogenoalkanes (using aqueous sodium hydroxide and ethanolic ammonia, respectively), and halogenoalkanes may be synthesised from alcohols using phosphorus halides, and how these are examples of moving within the level

Y

C

16

Halogenoalkanes

3 Substitution with ammonia, NH3 (in ethanol)

220

220

FO R

A3.3a

Original material © Cambridge University Press 2016

Location

Coursebook

(C)oursebook or (S)upplement or (B)oth

Chapter number

Chapter title

A3.3b

(b) explain how alcohols and amines may be synthesised from halogenoalkanes (using aqueous sodium hydroxide and ethanolic ammonia, respectively), and halogenoalkanes may be synthesised from alcohols using phosphorus halides, and how these are examples of moving within the level

A3.3c

Y

C

17

Alcohols, esters and carboxylic acids

(c) understand how substitution of halogenoalkanes by cyanide brings the carbon down from the Alcohol level to the Hydrocarbon level

Y

B

16

Halogenoalkanes

A3.3d

(d) understand how oxidation, including aldehydes and ketones from primary and secondary alcohols using acidified dichromate(VI), is an example of moving a carbon atom up a level.

Y

B

17

Alcohols, esters and carboxylic acids

A3.4

A3.4 Lower functional group level reactions – Carbonyl level Content

– addition of bisulfite • moving up a level: – oxidation reactions • moving down a level: – addition reactions students should be able to:

EV IE

• moving within the level: – hydrolysis to aldehyde or ketone

Supplement

Section number

Section title

Start page

End page

To add to which chapter of Coursebook?

Section number

Title

N

Present?

2 Substitution to form a halogenoalkane

227

228

2 Substitution with cyanide ions, CN – (in ethanol)

219

220

S17

S17.3

Halogenoalkane and cyanide ion

6 Oxidation

230

231

S17

S17.4

Oxidation increases the level

O

Syllabus content

W

Section number

LY

Cambridge International Pre-University Chemistry

(a) appreciate that hydrolysis of any member within the group leads to an aldehyde or ketone, and that this is an example of moving within the level

Y

B

18

Carbonyl compounds

The homologous series of aldehydes and ketones

235

236

S18

S18.1

Moving within the functional group level – hydrolysis

A3.4b

(b) understand that addition (where the π bond of the carbonyl breaks) of bisulfite (hydrogensulfate(IV)) is an example of reaction within the level

Y

B

18

Carbonyl compounds

Testing for the carbonyl group

238

239

S18

S18.2

Moving within the level – addition

A3.4c

(c) understand and recall that oxidation, including the use of acidified dichromate(VI) and Tollens’ reagent to form carboxylic acids from aldehydes, is an example of moving up a level

Y

B

18

Carbonyl compounds

Distinguishing between aldehydes and ketones

239

239

S18

S18.3

Moving up a level

A3.4d

(d) understand and recall that addition of hydrogen cyanide, Grignard reagents or metallic hydrides (e.g. NaBH4), where the π bond of the carbonyl breaks and a new C–C or C–H bond forms, is an example of a carbon atom moving down a level.

Y

B

18

Carbonyl compounds

Nucleophilic addition with HC N

237

238

S18

S18.4

Moving down a level

FO R

A3.5

R

A3.4a

A3.5 Addition and elimination reactions Content

• C=C and C=O in terms of dipole moments • addition reactions to C=C

Original material © Cambridge University Press 2016

Syllabus content

Present?

Location

Coursebook

(C)oursebook or (S)upplement or (B)oth

Chapter number

Chapter title

Supplement

Section number

• addition polymerisation • elimination reactions

students should be able to: A3.5a

(a) compare C=C and C=O π bonds in terms of dipole moments to explain why there is no nucleophilic attack on C=C

Y

B

18

Carbonyl compounds

A3.5b

(b) describe addition reactions to C=C: the reaction with H3O +, HBr, H2 and Br2

Y

B

15

Hydrocarbons

A3.5c

(c) describe addition polymerisation as an example of an addition reaction and in terms of the repeat unit in the polymer; suggest the monomer given the structure of an addition polymer

Y

C

15

A3.5d

(d) understand that elimination reactions are essentially the reverse of addition reactions and that elimination competes with substitution, the former being favoured by high temperature and high pH

Y

B

16

A3.5e

(e) describe the formation of alkenes by the elimination of HX from alkyl halides or H2O from alcohols.

Y

C

16

A3.5e

(e) describe the formation of alkenes by the elimination of HX from alkyl halides or H2O from alcohols.

Y

C

17

A3.6

A3.6 Green chemistry

Start page

End page

To add to which chapter of Coursebook?

Section number

Title

• atom economy • reducing environmental impact

238

S18

S18.5

Nucleophilic addition to C=C

Addition reactions of the alkenes

208

209

S15

S15.5

Addition of aqueous acid, H3O +(aq)

Hydrocarbons

Addition polymerisation; Tackling questions on addition polymers

211

213

Halogenoalkanes

Elimination reactions

222

222

S16

S16.1

Elimination reactions

Halogenoalkanes

Elimination reactions

222

222

Alcohols, esters and carboxylic acids

5 Dehydration

230

230

W

238

R

Note that the application of green chemistry principles in question papers may not be restricted to organic chemistry.

Mechanism of nucleophilic addition

EV IE

Content

O

• formation of alkenes by elimination

students should be able to:

Section title

N

Section number

LY

Cambridge International Pre-University Chemistry

(a) use the concept of atom economy as a measure of the efficiency of use of reagents in a synthesis (expressed as a percentage using [formula weight of utilised product(s)]/ [formula weight of all the reactants used] × 100%)

Y

S

S33

S33.3

Green Chemistry

A3.6b

(b) recall and discuss measures that can reduce the impact of chemical industry and research on the environment, including: finding benign alternatives to hazardous chemicals, using renewable feedstocks, using catalysts rather than stoichiometric reagents, etc.

Y

S

S33

S33.2, S33.4

Green Chemistry

FO R

A3.6a

Original material © Cambridge University Press 2016