The singlet in the spectrum occurs at δ 2.0 ppm, so this molecule is the one on the left in the diagram above –
ethyl ethanoate:
δ = 1.3 ppm δ = 4.1 ppm
H
H Hquartet O H δ = 2.0 ppm
C HC C O C C H singlet
triplet H
H
δ = 0.9–1.7 ppm H OH
O H CC
C C RO
RO
δ = 2.0–2.5 ppm
δ = 3.7–4.8 ppm
Notes on NMR spectra H
O Hb Hc
When a molecule is non-symmetrical and there are non-equivalent Ha C C C Hc
protons on both sides of a particular group, the spectrum becomes Ha Hb Hc
complex. Figure 11.32 shows the structural formula of propan-1-ol. In Figure 11.32 Propan-1-ol.
the NMR spectrum of this compound the Hb signal is split by Ha and
Hc. Depending on the strength of the coupling to each set of protons, CH3
the signal for Hb could either be described as a quartet of triplets or a HH
triplet of quartets.This signal is very complicated, and it is difficult to see
the exact nature of the splitting. One way of describing this peak is as a HH
complex multiplet. H
For a molecule such as methylbenzene (Figure 11.33), the protons Figure 11.33 Methylbenzene.
on the ring are not all equivalent, but because they are in very similar
environments they could show up as just one peak in the NMR spectrum δ = 3.4 ppm δ = 1.4 ppm
– unless a very high-resolution spectrum is generated. HHHH
δ = 0.9 ppm
The presence of highly electronegative atoms in a molecule (Figure
11.34) can cause chemical shifts to move to higher values.The closer the CI C C C C H
protons are to the very electronegative atom, the greater the effect. HHHH
δ = 1.7 ppm
Figure 11.34 Chlorine is more
electronegative than hydrogen and carbon.
11 MEASUREMENT AND DATA PROCESSING 543
? Test yourself 20 The NMR spectrum shown below is of an
18 Suggest the splitting pattern for each of the
following: alcohol. 6
a 1,1-dibromo-2,2-dichloroethane
b 1,1,3,3-tetrachloropropane
c propanoic acid
d HH H
HC COC H 1
1
HH H
4 32 1
19 The NMR spectrum of each of the following Chemical shift (δ)/ppm
contains a singlet. Suggest the chemical shift range Work out the structure of the alcohol and state the
multiplicity of the peak at δ = 4.0 ppm
for the singlet.
a propan-2-ol
b HO CH3 H
HCCOC CH
H HH
c
H OH
H C O C C CH3
HH
Single crystal X-ray crystallography
Further evidence for the structures of molecules can be obtained from
single crystal X-ray crystallography.This involves irradiating a crystal
with X-rays and looking at the positions and intensities of the diffracted
beams.This is an extremely powerful technique and gives a three-
dimensional picture of the molecule with bond lengths and bond angles
(Figure 11.35). If a single crystal can be grown, X-ray crystallography
usually provides the final word on the structure.
O(62) O(53) O(111) O(113)
O(11)
O(64)
O(112) O(13)
Os(6) P(11)
O(61) Os(5) O(31) Os(1) O(21)
O(63) P(1) O(12) O(24)
O(52) O(41) Os(4) Os(3) Os(2)
O(42) O(51) O(22) O(23)
O(33)
O(43) O(32)
Figure 11.35 The structure of a molecule obtained by X-ray crystallography.
544
Relative abundance / %Using combined spectroscopic techniques to
determine the structure of a molecule
In practice, information from more than one spectroscopic technique is
often used to determine the structure of a molecule.This is best illustrated
using some worked examples.
Worked examples
11.8 The mass spectrum below is for a compound, X, that contains carbon, hydrogen and oxygen.The infrared
spectrum of X has an absorption band at 1725 cm−1.
100
B
80
A
60
40
20
0 40 7780 105 120 136 160
0 m/z
a What is the relative molecular mass of X?
b Given that X contains two oxygen atoms, suggest the molecular formula for X.
c Identify which group is lost to form the peak marked B.
d Suggest the identity of the species responsible for the peak marked A.
e Calculate the IHD for X.
f Deduce the structural formula of X.
a Using the mass spectrum, the peak at highest m/z value represents the molecular ion – so the relative molecular
mass is 136.
b The only possible molecular formula containing two oxygen atoms that adds up to 136 is C8H8O2.
Two oxygen atoms have a combined mass of 32 and if we subtract this from 136 we get 104.The only other
elements in the molecule are carbon and hydrogen, so we must make 104 from a combination of the masses of
these. Seven carbon atoms gives a combined mass of 7 × 12 = 84, which would require there to be
104 − 84 = 20 hydrogen atoms in the molecule. However, the alkane with seven carbon atoms has only 16
hydrogen atoms, therefore C7H20O2 is not possible
c Peak B occurs at m/z 105. Now, 136 − 105 = 31. So we must think of combinations of C, H and O that add up
to 31.The only possible combination is CH3O, so this group is lost from the molecular ion to produce peak B.
(Note that C2H7 also adds up to 31, but the maximum number of hydrogen atoms for two carbon atoms is six.)
d Peak A occurs at m/z 77.The fragment responsible for this is C6H5+.
11 MEASUREMENT AND DATA PROCESSING 545
e IHD = 1 × [2c + 2 – h – x + n] for a molecule with formula CcHhNnOoXx, where X is a halogen atom.
2
The molecular formula is C8H8O2, so c = 8 and h = 8.
IHD = 1 × [(2 × 8) + 2 – 8] = 5
2
f A benzene ring has an IHD of 4 because it has a ring and three double bonds – so we can conclude that, apart
from the benzene ring, there is one other double bond or ring present. OH
The peak at m/z 77 indicates that the compound contains a benzene ring substituted
in one position.The loss of 31 mass units from the group suggests that it contains the CO C H
–O–CH3 group.The presence of an absorption band at 1725 cm−1 in the infrared H
spectrum suggests (using Table 11.11) that the molecule contains a C=O group.The
only possible molecule that contains all these elements is:
11.9 An organic compound, Y, is shown by elemental analysis to contain only carbon, hydrogen and oxygen.The
following information about the structure of Y is available:
• mass spectrum: the molecular ion peak occurs at m/z 86
• IHD = 1
• infrared spectrum:
100
% tranmittance
0 3000 2000 1000
• nuclear magnetic resonance spectrum: shows two peaks – one quartet (δ 2.4ppm) and one triplet
(δ 1.1ppm)
Determine the structure of the molecule from these data.
Adding up possible combinations of C(12), H(1) and O(16) to make the molecular ion peak at m/z 86
produces C5H10O, C4H6O2, C3H2O3 as possible formulas.
The IHD can be worked out for each of these molecular formulas using
1
IHD = 2 × [2c + 2 − h − x + n] for a molecule with formula CcHhNnOoXx, where X is a halogen atom.The
IHD corresponding to the possible formulas are:
C5H10O: 1 × [(2 × 5) + 2 − 10] = 1
C4H6O2: 12 × [(2 × 4) + 2 − 6] = 2
C3H2O3: 21 × [(2 × 3) + 2 − 2] = 3
2
546
Because the IHD is given as 1, the molecular formula must be C5H10O.
The infrared spectrum has an absorption band in the range 1700–1750 cm−1 indicating a C=O group in an
aldehyde, ketone, carboxylic acid or ester. Because there is only one oxygen atom in the molecule, we can rule
out esters and carboxylic acids because they both have two oxygen atoms in the functional group.Therefore the
compound must be an aldehyde or a ketone.
The NMR spectrum indicates that there are only two different chemical environments for the hydrogen atoms
in the molecule.There are several different isomers that can be drawn for a compound with molecular formula
C5H10O containing a C=O group – but only the two below have just two different chemical environments for
protons:
HHOHH CH3 O
H C C C C C H and H3C C C H
HH HH CH3
The splitting pattern from the NMR data can be used to distinguish between these two structures:
• a quartet signal indicates the presence of hydrogen atoms adjacent to a C atom with three H atoms attached
• a triplet signal indicates the presence of hydrogen atoms adjacent to a C atom with two H atoms attached.
These two signals together indicate the presence of an ethyl group.This group is present in only the left-hand
structure – so we can conclude that Y is pentan-3-one:
HHOHH
HCCCCCH
HH HH
The NMR spectrum of the right-hand molecule would consist of two singlets because all the hydrogen atoms
have no hydrogen atoms on the adjacent carbon.
As a final check you should use Table 11.14 to compare the chemical shift values given with what you would
expect for pentan-3-one.
Nature of science
Advances in technology have allowed scientists to derive ever more
detailed information about the structure of compounds.Theories of
structure and bonding have developed further as more information has
become available.
11 MEASUREMENT AND DATA PROCESSING 547
Exam-style questions
1 Rosie carried out an experiment in which she measured a temperature change. Her data are shown in the table.
Initial temperature / °C 18.7 ± 0.5
Maximum temperature / °C 37.6 ± 0.5
The temperature change should be quoted as: C 18.9 ± 1.0 °C
D 19.0 ± 1.0 °C
A 18.9 ± 0.5
B 19 ± 1 °C
2 Aazaish obtained the value 0.002 560 m3 from an experiment.The number of significant figures and decimal places is:
Significant figures Decimal places
A4 6
B6 4
C6 6
D3 6
3 Molly carried out an experiment to measure the enthalpy change of solution of a salt. In order to calculate a final
value, the following calculation was carried out:
[(50 ± 1) × 4.2 × (20 ± 1)]
1000 × (0.10 ± 0.01)
Quantities without uncertainties can be assumed to be exact. How should the final value be quoted?
A −42 ± 7 kJ mol−1 C −42.00 ± 2.01 kJ mol−1
B −42 ± 2 kJ mol−1 D −42.1 ± 7.1 kJ mol−1
4 Which of the following would be a good method for reducing the random uncertainty in an experiment to
measure the enthalpy change of neutralisation when 50 cm3 of 0.50 mol dm−3 sodium hydroxide reacts with 50 cm3
of 0.50 mol dm−3 hydrochloric acid?
A Insulate the reaction vessel with cotton wool.
B Stir the mixture more rapidly.
C Repeat the experiment.
D Measure out the liquids using a 50 cm3 measuring cylinder instead of a burette.
548
5 The graph shows the results of a series of experiments to investigate how the rate of the reaction X → Y varies
with the concentration of X.
Rate of reaction / mol dm–3 s–1 1.4
1.2
0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8
1 Concentration / mol dm–3
0.8
0.6
0.4
0.2
0
0
The gradient (slope) of this graph is: C 2.0 s−1
D 0.5 mol−1 dm3
A 2.0 mol dm−3 s−1
B 0.5 s
6 Which of the following would produce only a single peak in the low-resolution 1H NMR spectrum?
A CH3CH2OH C CH3CHO
B (CH3)3CCl D CH2ClCHClCH2Cl
7 Which of the following has an index of hydrogen deficiency of 2?
A C3H6O2 C C4H7N
B C6H6 D C5H8Cl2
8 Which of the following is unlikely be a peak in the mass spectrum of propanoic acid?
A m/z 45 C m/z 74
B m/z 15 D m/z 50
HL 9 The following information is provided about a compound M:
• There is only one peak in the NMR spectrum.
• IHD = 1.
• The molecular ion peak occurs at m/z 56.
• There is only one major band above 1500 cm−1 in the IR spectrum.
Which of the following could be the structure of compound M?
A H2C CH2 C HHH H
H2C CH2 HCCC C H
HH
B HOH D NH2
HCCCH HC C N
HH
H
11 MEASUREMENT AND DATA PROCESSING 549
HL 10 Which of the molecules below will not have a triplet in its NMR spectrum?
A HH C HH O
HCC H HCC C
H Br HH OH
bromoethane
D OH
B H Br Br H
HH CCH
HCCCCH HCCO H
HHHH
HH
11 Two separate experimental methods were used to determine the value of a particular experimental quantity. Each
experiment was repeated five times.The values obtained from these experiments are shown in the table.
Experiment A B
Trial Value Value
1 49.7 50.6
2 53.2 51.2
3 51.5 51.1
4 52.3 50.8
5 49.2 51.0
The literature value for this quantity is 50.9. [2]
a Explain which set of experimental values is more precise. [3]
b Work out a mean value for each experiment and use this to explain which set of data is more accurate.
12 Yi Jia carried out an experiment to measure a certain quantity.The value she obtained was 56.1 ± 0.5 kJ.The
literature value for this quantity is 55.2 kJ.
a Calculate the percentage error for this experiment. [1]
b The student maintained that any errors could be explained solely by random uncertainties. [2]
Is she correct? Explain your answer.
13 Jerry carried out an experiment to measure the rate of the reaction between Time / s Volume / cm3
magnesium and hydrochloric acid. He did this by recording the volume of
hydrogen gas collected every 15 s. His data are shown in the table. 0 0
15 19
a Plot a graph of these data. [3] 30 33
45 44
b Use your graph to determine the initial rate of reaction – include units. [3] 60 50
75 54
90 56
105 57
120 57
550
14 Certain molecules absorb infrared (IR) radiation.
a IR spectroscopy is one of the techniques that is used to determine the structure of organic molecules.
A student recorded the IR spectrum of a compound, X:
100
% Transmittance
0 3000 2000 1000
3800 Wavenumber / cm–1
The student knew that X was one of the following compounds:
HHHH I
II
HO C C C C H III
IV
H CH3 H H
HHH O
HC C C C
H H CH3 H
HHH O
HC C C C
H H CH3 OH
HHHH OHHH
HC C C C OC C C C H
H H CH3 H CH3 H H
Deduce which of the molecules is X and explain your choice by reference to the spectrum. [4]
[4]
b Explain why NMR is more useful than IR spectroscopy in distinguishing between propanal
and propanone.
11 MEASUREMENT AND DATA PROCESSING 551
15 The NMR spectrum of propan-2-ol is shown below.
432 1
Chemical shift (δ)/ppm
a Draw the full structural formula of propan-2-ol. [1]
[2]
b Explain why there are three peaks in the NMR spectrum of propan-2-ol.
[2]
c The integration trace is shown on the spectrum. Explain what information can be obtained from the [1]
integration trace.
[2]
d i Draw the full structural formula of butan-2-ol.
ii State the number of peaks and the relative areas under the peaks in the NMR spectrum of
butan-2-ol.
HL 16 Compounds may be identified using a combination of spectroscopic techniques.The mass spectrum for a
compound, Z, is shown below. Elemental analysis showed that Z contains C, H and O.
100
80
Relative abundance % 60
40
20
0
0 20 30 40 50 60 70
m/z
a Determine the relative molecular mass of Z. [1]
b Work out two possible molecular formulas for Z. [2]
c Deduce the formula of the fragments responsible for the peaks at m/z 31 and m/z 29. [2]
d The IR spectrum of Z shows a strong absorption band at about 3350 cm−1 but no absorptions in the
range 1600–1800 cm−1. Explain what conclusions about the structure of Z can be drawn from these data. [2]
552
e On the basis of the information above, suggest two possible structural formulas for Z. [2]
f The low-resolution NMR spectrum of Z is shown below.
432 1
Chemical shift (δ)/ppm
Explain how information from this spectrum can be used to deduce the structural formula of Z [4]
and draw the structural formula of Z. [2]
HL 17 NMR spectroscopy is a very powerful tool in the identification of organic compounds. [4]
a Predict and explain the splitting pattern of the hydrogen atom marked bold in the molecule below.
H HH
HC C CO
H H H
H CH
H
b A compound with the formula C4H10O has the NMR spectrum shown below. Draw the structural
formula of the compound and explain your reasoning.
4 3 21
Chemical shift (δ)/ppm
11 MEASUREMENT AND DATA PROCESSING 553
c There are two esters that have the molecular formula C3H6O2. Draw the structural formulas of these [5]
esters and explain the differences between their NMR spectra. [6]
d There are four esters with the molecular formula C4H8O2. Draw the structural formulas of these
esters and explain which will give rise to the NMR spectrum shown below.
4 3 21
Chemical shift (δ)/ppm
554
Summary Random uncertainties in analogue instruments
experimental values are – uncertainty is half the
Measurements are distributed caused by limitations of the smallest division
either side of the mean. measuring apparatus and digital instruments – uncertainty
other uncontrollable factors. is equal to the smallest division
can never be completely
eliminated but their effect can add absolute add percentage
be reduced by repeating the uncertainties together uncertainties together
experiment.
No figures should be quoted after
the uncertainty.
The uncertainty is in the adding and subtracting multiplying and dividing
last significant figure of the give answer to give answers to
measured value. fewest number of fewest number of
decimal places significant figures
given to one significant figure UNCERTAINTIES calculations
precision – how close repeat accuracy – how close measured cannot be reduced
readings are to each other value is to actual value by repeat readings
possible to have great systematic error – introduced
precision but low accuracy due to apparatus or procedure
shifts measured value in existence indicated by
one direction away from percentage error > total random uncertainty
the actual value
percentage error = |experimental value – accepted value| × 100%
|accepted value|
11 MEASUREMENT AND DATA PROCESSING 555
Summary – continued
IHD gives the number of HL
double bond equivalents in Single crystal X-ray
a molecule. For CcHhNnOoXx crystallography IR spectroscopy Infrared radiation
1 + 2 – h – x + n] provides a 3-D picture
IHD = 2 × [2c of a molecule with bond
lengths and bond angles.
SPECTROSCOPIC allows identification of
TECHNIQUES bonds/functional groups
present in a molecule
mass spectrometry
HL fragmentation pattern molecular ion peak
chemical shift gives provides information gives the relative
information about about the structure of molecular mass
the environment that a molecule
radio frequencies used protons are in
proton NMR TMS used as a only positive ions
reference standard produce a peak in
to set the zero on the the mass spectrum
chemical shift scale
only H nuclei integration trace
produce signals gives information
about the ratio of
the numbers of H in HL multiplicity – 1 = number
each environment splitting pattern gives of H atoms on adjacent C
information about the
number of peaks gives number of H atoms on
the number of different adjacent atom
chemical environments
for H atoms/protons
556
Group
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
1.01 4.00
H He
Appendix: the periodic tableKey
hydrogen 9.01 10.81 12.01 14.01 16.00 19.00 helium
1557 relative atomic mass 2
Be B C N O F
6.94 atomic symbol 20.18
beryllium
Li 4 name Ne
atomic number
lithium 24.31 boron carbon nitrogen oxygen fluorine neon
3 5 6 7 8 9 10
Mg
22.99 26.98 28.09 30.97 32.07 35.45 39.95
magnesium
Na 12 Al Si P S Cl Ar
sodium 40.08 aluminium silicon phosphorus sulfur chlorine argon
11 13 14 15 16 17 18
Ca
39.10 44.96 47.87 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.38 69.72 72.63 74.92 78.96 79.90 83.80
calcium Ti V Cr Mn Fe Co Ni Cu Zn Ga As Se
K 20 Sc Ge Br Kr
potassium 87.62 scandium titanium vanadium chromium manganese iron cobalt nickel copper zinc gallium germanium arsenic selenium bromine krypton
19 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36
Sr
85.47 88.91 91.22 92.91 95.96 (98) 101.07 102.91 106.42 107.87 112.41 114.82 118.71 121.76 127.60 126.90 131.29
strontium Zr Nb
Rb 38 Y Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
rubidium 137.33 yttrium zirconium niobium molybdenum technetium ruthenium rhodium palladium silver cadmium indium tin antimony tellurium iodine xenon
37 39 40 41 42 43 44 48 49 50 54
Ba 45 46 47 51 52 53
132.91
barium 138.91 * 178.49 Ta180.95 W183.84 Re186.21 Os190.23 Ir192.22 Pt195.08 Au196.97 Hg200.59 Tl204.38 2P07b.2 Bi208.98 (P20o9) (A21t0) R(22n2)
Cs 56
La Hf
caesium (226)
55 lanthanum hafnium tantalum tungsten rhenium osmium iridium platinum gold mercury thallium lead bismuth polonium astatine radon
Ra 57 72 73 74 75 76 80 81 82
(223) 77 78 79 83 84 85 86
radium
Fr 88 (227) ** (267) (268) (269) (270) (269) (278) (281) (281) (285) (286) (289) (28) (293) (294) (294)
Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Fl Uup Lv Uus Uuo
francium Ac
87
actinium rutherfordium dubnium seaborgium bohrium hassium meitnerium darmstadtium roentgenium copernicium ununtrium flerovium ununpentium livermorium ununseptium ununoctium
104 105 106 107 108 109 110 111 112 113 114 115 116 117 118
89
lanthanoids * 140.12 140.91 144.24 145 150.36 151.96 157.25 158.93 162.50 164.93 167.26 168.93 173.05 174.97
actinoids **
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
cerium praseodymium neodymium promethium samarium europium gadolinium terbium dysprosium holmium erbium thulium ytterbium lutetium
58 59 60 61 62 63 64 65 66 67 68 69 70 71
232.04 231.04 238.03 (237) (244) (243) (247) (247) (251) (252) (257) (258) (259) (262)
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
thorium protactinium uranium neptunium plutonium americium curium berkelium californium einsteinium fermium mendelevium nobelium lawrencium
90 91 92 93 94 95 96 97 98 99 100 101 102 103
Answers to test yourself questions
Topic 1 5 Compound Molar mass / Mass / g Number of
moles / mol
1 a 2NO + O2 → 2NO2 g mol−1
b C3H8 + 5O2 → 3CO2 + 4H2O
c CaCO3 + 2HCl → CaCl2 + CO2 + H2O H2O 18.02 9.01 0.500
d C2H5OH + 3O2 → 2CO2 + 3H2O
e WO3 + 3H2 → W + 3H2O CO2 44.01 5.00 0.114
f 2H2O2 → O2 + 2H2O
g 4CrO3 → 2Cr2O3 + 3O2 H2S 34.09 3.41 0.100
h Al4C3 + 6H2O → 3CH4 + 2Al2O3
i 8HI + H2SO4 → H2S + 4H2O + 4I2 NH3 17.04 59.6 3.50
j 4PH3 + 8O2 → P4O10 + 6H2O
Q 28.6 1.00 0.0350
Z 51.61 0.0578 1.12 × 10−3
2 a compound f element Mg(NO3)2 148.33 1.75 0.0118
C3H7OH 60.11 2500 41.59
Fe2O3 159.70 9.07 × 10−3 5.68 × 10−5
b element g compound 6 a 2.99 × 10−23 g c 7.31 × 10−23 g
b 2.83 × 10−23 g c 9.03 × 1022
c compound h mixture c 2.17 × 1023
7 a 1.20 × 1024
d mixture i compound b 4.82 × 1023
e mixture j compound 8 a 2.41 × 1022
b 9.63 × 1023
3 a compound; gas c mixture; compound;
b element; solid element; gas
d mixture; elements; gas 9 a 0.8 mol c 0.12 mol
b 0.7 mol
4 Compound Relative molecular mass
SO2 64.07 10 a 34.72% c 61.24%
NH3 17.04 b 43.19%
C2H5OH 46.08
MgCl2 95.21 11 a 1.60 g c 5.64 g
Ca(NO3)2 164.10 b 2.50 g
CH3(CH2)5CH3 100.23
PCl5 208.22 12 a 2.00 g c 2.29 g
Mg3(PO4)2 262.87
Na2S2O3 158.12 b 1.67 g
CH3CH2CH2COOCH2CH3 116.18
13 CO2 CH HO C3H8 H2O PCl5 C6H5CH3
14 Empirical Relative Molecular
formula molecular mass formula
HO 34.02 H2O2
166.90 Cl2O6
ClO3 C6H12
CH2 84.18 B3N3H6
BNH2 80.52
15 a CH3 b C2H6
16 Cl2O
17 I2O5
18 C3H8O
19 Fe2O3
558
20 a 0.2 mol d 0.9 mol Topic 2
b 0.05 mol e 0.6 mol
c 0.01 mol f 3.6 × 10−3 mol 1 23982U 92 protons 146 neutrons 92 electrons
c 4.62 g 7353As 33 protons 42 neutrons 33 electrons
21 a 0.363 g d 2.97 g 8315Br 35 protons 46 neutrons 35 electrons
b 2.67 g c 65.1%
c HNO3 2 4200Ca2+ 20 protons 20 neutrons 18 electrons
22 a 80.0% d NaCl 15237I− 53 protons 74 neutrons 54 electrons
b 94.7% 14508Ce3+ 58 protons 82 neutrons 55 electrons
b 1.6 dm3
23 a H2SO4 d 0.0176 mol 3 The only element is 11H, which has one proton and
b H2O e 0.0110 mol
d 19.3 dm3 no neutrons.
24 0.15 g e 13.6 dm3
4 a D and L are isotopes; Q and M are isotopes
25 3.10 tonnes b 300 cm3 b D, X, L
26 a 200 cm3
b 573 cm3 5 52.06
27 a 0.0106 mol b 0.200 mol dm−3
b 0.0881 mol c 0.0108 mol 6 28.11
c 0.00441 mol
b 0.284 dm3 7 a 91% indium-115; 9% indium-113
28 a 2.27 dm3 b 63.85% gallium-69; 36.15% gallium-71
b 2270 dm3
c 6.13 dm3 8 a microwaves infrared radiation orange light
green light ultraviolet radiation
29 114 cm3 infrared radiation orange light
30 1910 cm3 b microwaves ultraviolet radiation
green light
31 0.4669 g
32 a 290 cm3 9 An electron in a hydrogen atom that has been
33 21.7 cm3 promoted to a higher energy level falls down to
energy level 1.As it does so, it gives out energy
34 196 °C in the form of a photon of light in the ultraviolet
region of the electromagnetic spectrum.
35 0.0655 mol
36 43.2 g mol−1 10 4
37 28.2 dm3
38 a 0.115 dm3 c
39 0.782 g
3
40 a 3.55 g
41 a 4.40 × 10−3 mol b
b 2.34 × 10−3 mol increasing energy 2
42 0.475 mol dm−3
a
43 a 1.25 g
44 42.8 cm3 1
45 56.8 cm3
a any transition to level 1
b any transition to level 2
c any transition to level 3
11 a 1s22s22p3
b 1s22s22p63s23p2
c 1s22s22p63s23p6
d 1s22s22p63s23p64s23d104p3
e 1s22s22p63s23p64s23d3
ANSWERS TO TEST YOURSELF QUESTIONS 559
12 a 2p2 6 a 2Rb(s) + 2H2O(l) → 2RbOH(aq) + H2(g)
b 2K(s) + Br2(l) → 2KBr(s)
2s2 c Cl2(aq) + 2KBr(aq) → 2KCl(aq) + Br2(aq)
d Na2O(s) + H2O(l) → 2NaOH(aq)
1s2 e SO3(g) + H2O(l) → H2SO4(aq)
b 3p3 7 a same c different
b different d same
3s2
2p6 8 a acidic c alkaline
2s2 b alkaline
1s2 9 a 1s22s22p63s23p63d8 c 1s22s22p63s23p63d2
b 1s22s22p63s23p63d6 d 1s22s22p63s23p63d3
c
3d5 10 a +2 b +2 c +2 d +3
4s1 e +3 f +3 g +6 h0
3p6
3s2 11 Cu2+, Fe2+ and Co3+
2p6
2s2 12 MnF3, CoF2 b [Fe(H2O)6]2+
13 a [Co(H2O)6]2+
1s2
Topic 4
1 a ionic c ionic e covalent
b covalent d covalent f ionic
m AgNO3
13 a 4.85 × 10−19 J; b 2.31 × 1015 Hz 2 a MgO g Li3N n NH4Cl
h Mg3(PO4)2 o Cu(NO3)2
14 X is in group 14; Z is in group 2; Q is in group 15. b BaSO4 i MgF2 p Rb2CO3
15 a 1s22s22p63s23p6 c Ca(OH)2 j K2SO4
d Na2O k (NH4)2CO3 NO
b 1s22s22p63s23p63d3 e SrS l Ag2S
c 1s22s22p63s23p63d7 NNH
d 1s22s22p63s23p64s23d104p6 f Al2O3 HH
Topic 3 3 Na < H < Br < Cl < O gF
4a
SH
H
1 a silicon d iodine b Cl hH
b tellurium e arsenic P Cl
c polonium Cl
2 a increase b decrease c Cl iH O O H
Cl C Cl
3 a Mg < Ca < Sr < Ba Cl j +
b Na+ < F− < O2− Cl
c Al3+ < Na+ < Na < K dO
d Cl < S < Cl− < S2− < I− CF
F
4 a false d false Cl P Cl
b false e true Cl
c true
k N O+
5 Element X eH C N l O C N–
fS C S
560
5a F eN N O 7 HBr polar δH+ Bδ–r
F Xe F or δH+ C Nδ–
F HCN polar
NNO δH+ δP– Hδ+
b Cl – fF PH3 polar Hδ+
Cl Cl
P FSF SCl2 polar Cδ–l δS+ Cδ–l
Cl Cl CF4 non-polar
Cl F N2 non-polar
OCl2 polar
cF g I– BCl3 non-polar Cδ+l Oδ– Cδ+l
FF I C2Cl2 non-polar
Br I H2S polar δH+ δS– δH+
FF
CH2Cl2 polar δH+
dF h N N N– Cδ–l C δH+
Clδ–
or
Cl F N N N– 8 a CH4 CF4 CCl4
F
b PH3 AsH3 NH3
c CH4 NH3 N2H4
6 a bent/V-shaped/angular; predict 100 –108° (actual d C2H4 CH3F CH3OH
value 92°)
e H2S H2O H2O2
b trigonal pyramidal; predict 100 –108° (actual
value 100°) f CH3CH2CH2CH2CH3 CH3CH2OCH2CH3
c tetrahedral; 109.5° CH3CH2CH2CH2OH
d linear; 180°
e trigonal planar; predict F–C–F bond angle of g Ne N2 F2 HF
about 118–120° (greater repulsion due to double 9 a HCl CCl4 SiCl4 NaCl
bond) and O–C–F bond angle of about 120– PBr3
122° (actual values: F–C–F bond angle = 108°, b HBr Br2 C3H7OH CaBr2
O–C–F bond angle = 126°) C4H9OH
f linear; 180° c C3H8 C4H10 SiO2
g bent/V-shaped/angular; predict about 118°
(based on trigonal planar) (actual value 110°) CH3CH2COOH
h tetrahedral; 109.5°
i linear; 180° d CO2 SO2
j bent/V-shaped/angular; predict about 118°
(based on trigonal planar) (actual value 117°) 10 a C6H12 C5H11OH NaCl
k trigonal planar about each C; predict F–C–F b CH4 CH3Cl CaCl2
bond angle of about 118° and F–C–C bond angle
of about 122° (actual value for F–C–C bond 11 a K Na Li
angle about 124°)
l linear; 180° b Na Mg Al
12 a Lewis structures: Formal charges:
OO O O 1−
Cl P Cl Cl P Cl Cl P Cl Cl P 1+ Cl
Cl
Cl Cl Cl
preferred
ANSWERS TO TEST YOURSELF QUESTIONS 561
b Lewis structures: Formal charges: BrF3 polar Fδ–
Fδ– Brδ+
O– O–O O 1−
Fδ–
O Cl O O Cl O O Cl 1− 1− Cl O3+ 1−
OO
O O O O 1− SOCl2 polar Clδ– Sδ+ Oδ–
Clδ–
preferred
c Lewis structures: Formal charges: 15 a 1σ 1π e 3σ 1π i 8σ 1π
b 1σ 2π f 2σ 2π j 3σ 1π
OO O O 1− c 3σ g 3σ 1π k 6σ 2π
d 2σ 2π h 1σ 2π l 6σ 2π
O Xe O O Xe O O Xe O O 1− Xe4+ O 1−
OO O O 1− 16 a 1.5
preferred b 1.25
13 a T-shaped (arrow-shaped); predict bond angles of c 1.75
about 88° (actual value 86°)
17 a sp2 e sp i sp2
b square-based pyramid; predict bond angle of b sp3 f sp3 j sp2
about 88° c sp3 g sp3
d sp3 h sp
c (distorted) tetrahedral; predict bond angles
O–S–O of about 112° and Cl–S–Cl of about 18 a sp3 b sp2
107° (angles of 109° are fine!) (actual values:
O–S–O = 120° and Cl–S–Cl = 111°) 19 a sp2 c sp3
d sp2
d see-saw shaped; predict bond angles of about 88° b sp
and 118° (actual values 85° and 101°)
Topic 5
e linear; 180°
f octahedral; 90° 1 a 0.128 J g−1 °C−1 b 0.236 J g−1 °C−1
g square-based pyramid; predict bond angle of
2 a −4200 kJ mol−1 b −5490 kJ mol−1
about 88° (actual value 79°)
h see-saw shaped; predict bond angles of about 88° 3 −1080 kJ mol−1
and 118° 4 Heat energy loss to the surroundings; incomplete
i bent/V-shaped/angular; predict bond angles of
combustion (other, more minor, factors include
about 105° (actual value 102°)
j linear; 180° evaporation of water and/or propan-1-ol)
5 a −56.8 kJ mol−1 b i 1.36 °C
ii 2.72 °C
14 XeF6 non-polar iii 2.72 °C
XeF4 non-polar
SF4 polar Fδ– 6 a −151 kJ mol−1 c 15.2 °C
Fδ– b 7.6 °C
Sδ+ 7 a +44 kJ mol−1 b +31 kJ mol−1
δ–F
8 a −400 kJ mol−1 b −91 kJ mol−1
Fδ–
PCl5 non-polar δF– δS+ δF– 9 −312 kJ mol−1
SF2 polar
10 –126 kJ mol−1
SF6 non-polar 11 a 1 H2(g) + 1 F2(g) → HF(g)
ClF5 polar 2 2
3 12Cl2(g)
b 2 H2(g) + C(s) + → CH3Cl(g)
Fδ– c H2(g) + 21O2(g) → H2O(l)
δ–F Cδl + Fδ– 1
δ–F Fδ– d 6H2(g) + 5C(s) + 2 O2(g) → C5H11OH(l)
562
12 +20 kJ mol−1 Topic 6
13 −574.1 kJ mol−1
1a
14 +33 kJ mol−1
15 a −115 kJ mol−1 c −107 kJ mol−1 80
b −287 kJ mol−1 70
60
16 147 kJ mol−1 Volume of gas / cm3 50
40
17 124.5 kJ mol−1 30
20
18 −273 kJ mol−1 10
19 −344 kJ mol−1
0
20 2307 kJ mol−1 0 50 100 150 200
Time / s
21 KCl < LiF < CaCl2 < CaS < CaO
The gradient is steepest at the beginning – the
22 a always endothermic reaction is fastest; the concentration of reactants
b always endothermic is highest at the beginning, so the collision
c always endothermic frequency is highest.
d sometimes exothermic and sometimes The gradient decreases as the reaction goes on
endothermic – the reaction rate decreases; reactants are being
e always exothermic used up – there is a lower concentration of
f always endothermic reactants, so lower collision frequency.
The graph becomes horizontal – the reaction
23 14 kJ mol–1 has finished; the acid is in excess, so the reaction
finishes when all the magnesium has been used
24 –155 kJ mol–1 up.
b The initial rate is approximately 1.3 cm3 s−1.
25 a decrease c increase c The average rate is 0.62 cm3 s−1.
b decrease d increase
d, e, f
26 a −38 J K−1 mol−1 c 411 J K−1 mol−1
b 763 J K−1 mol−1
27 ∆G = −121 kJ mol−1
The reaction is spontaneous; ∆G is negative.
28 ∆G = −2108 kJ mol−1 80
70
The reaction is spontaneous; ∆G is negative.
X
29 a 889 J K−1 mol−1 b 399 °C 60 Z
30 a more spontaneous 50
Volume of gas / cm3 40 Y
b less spontaneous 30 200
c more spontaneous 20
d more spontaneous 10
31 Because ∆G is positive, the position of equilibrium 0 100 150
lies closer to the reactant (N2O4) 0 50 Time / s
g
lower temperature
Number of particles higher temperature
Ea greater number of
particles with energy
greater than or equal to
the activation energy
0
0 Energy
ANSWERS TO TEST YOURSELF QUESTIONS 563
At the higher temperature, there are more 9 X +Y RDS XY
particles with energy greater than the activation
energy.Therefore there is a greater chance that ⎯→
a collision will result in a reaction and more
successful collisions per unit time. XY + X fast 2Z
10 56 kJ mol−1; A = 1.97 × 108 mol−1 dm3 s−1
11 90 kJ mol−1; A = 9.74 × 109 mol−1 dm3 s−1
2 a First order with respect to A; second order with Topic 7
respect to B 1 a shifts to left c shifts to right
b shifts to left
b3
c 10 mol−2 dm6 s−1
d 2.50 × 10−3 mol dm−3 s−1
3 a rate = k[X][Y] 2 a shifts to left c shifts to left
b shifts to right
b k = 3.96
c mol−1 dm3 s−1 3 a shifts to right
b shifts to left
4 The rate of reaction would be quadrupled c shifts to right
(increased by a factor of 4). d shifts to right (the products formed are acids and
will react with NaOH – their concentration is
5 a First order lowered)
b Zero order 4 a Kc = [N2O3(g)]
[NO(g)][NO2(g)]
c rate = k[D]
d k = 0.171 s−1
e 0.0103 mol dm−3 s−1
6 • The mechanism does not agree with the b Kc = [CO(g)][H2(g)]3
stoichiometric equation (the overall equation [CH4(g)][H2O(g)]
from this mechanism is 2A + 3B → 3C + 2D).
c Kc = [H2(g)]2[O2(g)]
• The rate-determining step contains A, but A does [H2O(g)]2
not appear in the rate equation.
d Kc = [NO(g)]4[H2O(g)]6
• Step 2 involves three molecules colliding, which [NH3(g)]4[O2(g)]5
is extremely unlikely. [NO2(g)]2
[NO(g)]2[O2(g)]
7 a rate = k[Q]2 e Kc =
b Because it does not agree with the stoichiometric
equation – the overall equation from this 5 a √1033 or 3.2 × 1016
mechanism would be: 2Q + P + Z → Y + R + S
c Mechanism 3 b 10−19
c √10−19 or 3.2 ×10−10
fast 6 a 9.71 × 10–13 b 7.73
8 NO + NO N2O2 7 a shifts to left; Kc stays the same
N2O2 + Br2 RDS 2NOBr b shifts to right; Kc increases
⎯→
or c shifts to right; Kc increases
d shifts to right; Kc stays the same
NO + Br2 fast NOBr2
e shifts to right; Kc stays the same
RDS
NOBr2 + NO 2NOBr f no effect; Kc stays the same
⎯→
b 1.02 × 10−9
or 8 a 22.2
NO + Br2 fast NOBr + Br 9 a 0.520 mol Z
RDS b 0.600 mol A; 1.800 mol X
NO + Br ⎯→ NOBr
564
10 a Kc = [Q(g)] 9 a calcium oxide/hydroxide and nitric acid
[A(g)][X(g)] b cobalt(II) oxide/hydroxide and sulfuric acid
c copper(II) oxide/hydroxide and hydrochloric acid
b i 3.33 d magnesium oxide/hydroxide and ethanoic acid
ii 200 10 0.10 mol dm−3 H2SO4 < 0.10 mol dm−3 HCl
< 0.010 mol dm−3 HCl < 1.0 mol dm−3 NaOH
c endothermic
11 a false b true c false d false
11 a Kc = [Q(g)]4[Z(g)]
[A(g)]2[X(g)] 12 a X = 1.7; Y = 3.7
b measure electrical conductivity: X higher; add
b i 2.37 × 10−4 magnesium or calcium carbonate: X will react
more vigorously
ii 0.154
c exothermic
12 a 0.075 mol X 13 a solution containing 0.001 00 mol dm−3 H+(aq)
b 0.080 mol dm−3 of A
13 a –2.53 kJ mol–1 3
b –35.4 kJ mol–1
c 43.5 kJ mol–1 a solution containing 1.00 × 10−12 mol dm−3 H+(aq) 12
a solution of 1.00 mol dm−3 HCl(aq) 0
14 a 1.70 × 10–9 a solution of 2.00 × 10−4 mol dm−3 HNO3(aq) 3.70
b 5.79 × 103
c 6.72 × 1026 a solution of CH3COOH of concentration 2.30
0.100 mol dm−3 assuming 5% dissociation of the
acid
Topic 8 14 a 3.16 × 10−4 mol dm−3
b 1.26 × 10−8 mol dm−3
1 An acid is a proton (H+) donor; a base is a proton c 1.58 × 10−13 mol dm−3
(H+) acceptor.
15
2 CH3COOH(aq) + NH3(aq) [H+(aq)] / [OH−(aq)] / mol dm−3 pH Acidic or
acid 1 base 2 mol dm−3 alkaline?
1 × 10−11
CH3COO−(aq) + NH4+(aq) 1 × 10−3 1 × 10−9 3 acidic
base 1 acid 2 1 × 10−5 0.01 5 acidic
1 × 10−12 12 basic
3 a NH4+ c H2SO4 e H2PO4−
b H2O d HCN
16 a 12.2 b 13.0
4 a CO32− c HCOO−
b OH− d NH2− 17 a 5.0 × 10−13 mol dm−3
b 7.1 × 10−11 mol dm−3
5 a acid
b base 18 Strong Weak Strong Weak Salt
c base acid base base
acid NaOH NH4NO3
6 An acid is an electron pair acceptor; a base is an HCl H2CO3 Ba(OH)2 NH3 Na2SO4
electron pair donor. H2SO4 HCOOH CH3NH2 KNO3
HNO3
7 H2O = base; BF3 = acid; HCO3−= base; H+ = acid;
AlCl3 = acid; NH3 = base; CO = base 19 Hydrochloric acid is a strong acid and ethanoic acid
is a weak acid, but HCl has the higher pH.The pH
8 a Zn + H2SO4 → ZnSO4 + H2 is just a measure of the concentration of H+ ions in
b CuO + 2HNO3 → Cu(NO3)2 + H2O solution and therefore depends on concentration.
c 2NH3 + H2SO4 → (NH4)2SO4 pH is only a useful measure of acid strength
d Ca(HCO3)2 + 2HCl → CaCl2 + 2CO2 + 2H2O when solutions of equal concentration are being
e Mg(OH)2 + H2SO4 → MgSO4 + 2H2O compared.
f Cu + H2SO4 → no reaction
g CaO + 2HCl → CaCl2 + H2O
ANSWERS TO TEST YOURSELF QUESTIONS 565
20 a 0.10 mol dm−3 HCl 31
b 0.10 mol dm−3 KOH
c 0.10 mol dm−3 H2SO4 Acid Ka pKa Conjugate Kb pKb
d 0.10 mol dm−3 HNO3 base
HCN 3.98 × 10−10 9.40 4.60
21 Because carbon dioxide in the atmosphere dissolves HF 5.62 × 10−4 3.25 CN− 2.51 × 10−5 10.75
in rain to form carbonic acid, a weak acid. HIO3 0.158 0.8 13.2
NH4+ 5.62 × 10−10 9.25 F− 1.78 × 10−11
22 H2SO3, H2SO4, HNO2, HNO3, (H2CO3) CH3COOH 1.74 × 10−5 4.76 4.75
CH3NH3+ 2.29 × 10−11 10.64 IO3− 6.31 × 10−14 9.24
23 a NH3 1.78 × 10−5 3.36
CH3COO− 5.75 × 10−10
Acid Concentration [H+] / pH Ka pKa CH3NH2 4.37 × 10−4
of acid / mol dm−3 9.40 32 Acid
mol dm−3 8.70
6.90 Alkali Indicator
8.74
HA 0.0100 2.00 × 10−6 5.70 4.00 × 10−10 7.43 0.100 mol dm−3 0.100 mol dm−3 NaOH phenolphthalein
HB 0.200 2.00 × 10−5 4.70 2.00 × 10−9 5.30 CH3COOH
HC 0.500 2.50 × 10−4 3.60 1.25 × 10−7 0.020 mol dm−3 KOH bromothymol
HD 2.20 × 10−2 6.31 × 10−6 5.20 1.81 × 10−9 0.010 mol dm−3 0.010 mol dm−3 NH3 blue
HE 0.250 9.64 × 10−5 4.02 3.72 × 10−8 HNO3
HF 0.0300 3.86 × 10−4 3.41 4.96 × 10−6 bromocresol
0.010 mol dm−3 HCl green
b HF > HC > HE > HB > HD > HA 33 initial pH 3.23
24 a volume of NaOH required to reach the 50.0 cm3
equivalence point
Concentration [OH−] /
Base Kb pKb approximate pH at equivalence point 8–9 (8.3)
of base / mol dm−3 approximate final pH (after approaches 12
approximately 200 cm3 of NaOH has
mol dm−3 been added)
B1 0.100 1.33 × 10−3 1.77 × 10−5 4.75 34 4.80
B2 0.250 3.79 × 10−3 5.75 × 10−5 4.24 35
B3 0.0200 4.70 × 10−4 1.10 × 10−5 4.96
0.100 mol dm−3
b B3 < B1 < B2 CH3CH2CH2CH2COONa pH
0.500 mol dm−3 KNO3 greater than 7
25 HOI < HOCl < HNO2 < HClO2 0.100 mol dm−3 Na2CO3
0.100 mol dm−3 CH3CH2NH3+Cl− 7
26 a 3.28 b 4.85 c 6.30 0.200 mol dm−3 CrCl3 greater than 7
less than 7
27 a 7.37 b 6.84 c 6.14 less than 7
28 a 1.52 b 0.301 c 12.3
29 a 11.7 b 8.31
30 a 12.53 36 a 11.2 b 9.09
b 1.000 (pKw makes no difference to the 37 a yes c yes e no g yes
calculation because we ignore any dissociation of b no d yes f yes
water)
c 12.13
566
Topic 9 d N2 + 6H2O → 2NO3− + 12H+ + 10e−
e SO42− + 4H+ + 2e− → H2SO3 + H2O
1 a +3 c +1 e −2
b +2 d +4 f +3 11 a VO2+ + 2H+ + e− → V3+ + H2O
b Xe + 3H2O → XeO3 + 6H+ + 6e−
2 a +1 c +7 e +5 c NO3− + 4H+ + 3e− → NO + 2H2O
b −1 d +3 d 2NO3− + 10H+ + 8e− → N2O + 5H2O
e VO2+ + 2H+ + e− → VO2+ + H2O
3 a +1 c +4 e −3 (NH4+ NO3−)
b +5 d +6 f –1 12 a Cl2 + 2Br− → 2Cl− + Br2
b Zn + 2Ag+ → Zn2+ + 2Ag
4 a nitrogen(II) oxide d potassium iodate(V) c 2Fe3+ + 2I− → 2Fe2+ + I2
b chlorine(VII) oxide e chromium(III) oxide
c selenium(IV) oxide 13 a 6Fe2+ + Cr2O72− + 14H+ → 6Fe3+ + 2Cr3+ + 7H2O
b 6I− + Cr2O72− + 14H+ → 3I2 + 2Cr3+ + 7H2O
5 a +2 c +2 e +1 c Zn + 2VO2+ + 4H+ → Zn2+ + 2V3+ + 2H2O
b +3 d +2 f +2 d 2BrO3− + 10I− + 12H+ → Br2 + 5I2 + 6H2O
e 2H2O + NpO22+ + U4+ → NpO2+ + UO2+ + 4H+
6 a reduction d oxidation
b reduction e reduction 14 a 5.00 × 10−5 mol
c oxidation b 2.50 × 10−5 mol
c 2.50 × 10−5 mol
7 a not redox d 8.00 mg dm−3
b redox: S oxidised; O reduced e 3.40 mg dm−3; <5 ppm therefore not polluted
c redox: Na oxidised; H reduced
d not redox 15 a X > Z > Q
e not redox
f redox: Fe oxidised; O reduced (oxidation number b i stronger
of O in H2O2 is −1)
g redox: O oxidised; Hg reduced; S reduced ii X > Z > A > Q
h redox: I oxidised; Cl reduced iii Q2+
8 a CuSO4 is the oxidising agent; Zn is the reducing c i from X to A
agent
ii X
b Cl2 is the oxidising agent; Br− is the reducing agent iii X(s)|X2+(aq)||A2+(aq)|A(s)
c I2O5 is the oxidising agent; CO is the reducing
iv lower
agent
16 a anode: bromine cathode: potassium
d HNO3 is the oxidising agent; S is the reducing b anode: chlorine cathode: copper
agent c anode: oxygen cathode: nickel
d anode: chlorine cathode: calcium
e I2 is the oxidising agent; Na2S2O3 is the reducing
agent 17 a anode: 2Cl− → Cl2 + 2e−
cathode: Na+ + e− → Na
f KMnO4 is the oxidising agent; Na2C2O4 is the
reducing agent b anode: 2O2− → O2 + 4e−
cathode: Fe3+ + 3e− → Fe
g K2Cr2O7 is the oxidising agent; FeSO4 is the
reducing agent c anode: 2Br− → Br2 + 2e−
cathode: Mg2+ + 2e− → Mg
9 a Fe3+ + 3e− → Fe 18 a +0.21V; direction of electron flow from iron half-
b Pb2+ → Pb4+ + 2e−
c I2 + 2e− → 2I− cell to nickel half-cell
d 2S2O32− → S4O62− + 2e− Ni2+(aq) + Fe(s) → Ni(s) + Fe2+(aq)
e C2O42− → 2CO2 + 2e−
b +0.82V; from iodine half-cell to chlorine half-cell
10 a I2 + 2H2O → 2OI− + 4H+ + 2e− Cl2(g) + 2I−(aq) → 2Cl−(aq) + I2(aq)
b MnO4− + 4H+ + 3e− → MnO2 + 2H2O
c 2IO3− + 12H+ + 10e− → I2 + 6H2O c +1.56 V; from zinc half-cell to silver half-cell
Zn(s) + 2Ag+(aq) → Zn2+(aq) + 2Ag(s)
ANSWERS TO TEST YOURSELF QUESTIONS 567
d +0.56V; from iron half-cell to chromium half-cell 24 a 0.20 mol Na and 0.10 mol Cl2
6Fe2+(aq) + Cr2O72−(aq) + 14H+(aq) b 0.10 mol Cu and 0.10 mol Cl2
→ 6Fe3+(aq) + 2Cr3+(aq) + 7H2O(l) c 0.10 mol H2 and 0.050 mol O2
e +0.15 V; from chlorine half-cell to manganese Topic 10
half-cell 1 C15H32 Empirical formula
MnO4−(aq) + 16H+(aq) + 10Cl−(aq)
2 Molecular formula C3H7
→ 2Mn2+(aq) + 8H2O(l) + 5Cl2(g) CH2
f +1.53V; from iron half-cell to bromine half-cell a C6H14 C8H8O
b C5H10 C5H6O
Fe(s) + Br2(l) → Fe2+(aq) + 2Br−(aq) c C8H8O
d C10H12O2
19 a cell potential negative, therefore not spontaneous
3 a alcohol / hydroxyl / –OH group
b cell potential +0.24 V, therefore spontaneous;
Cr2O72− is the oxidising agent and Br− is the b alkene / C=C / alkenyl; carboxylic acid / carboxyl /
reducing agent COOH group
c cell potential negative, therefore not spontaneous c ketone / carbonyl
d ester
d cell potential negative, therefore not spontaneous e ether; nitrile
f benzene ring / phenyl; carboxylic acid / carboxyl
20 a i Po2+ ii Np g aldehyde / carbonyl; alkyne / alkynyl
b i true v false
ii true vi true
iii false vii true
iv true
c i not spontaneous
ii spontaneous
iii not spontaneous (both oxidation reactions)
21 a E cell = +0.53V; 4 a 2-methylpentane
ΔG = −100 kJ mol−1;
spontaneous b 2,2,3-trimethylbutane
c 2,2-dimethylpentane
b E cell = −0.47V; d 2-methylpentane
ΔG = +91 kJ mol−1; e but-2-ene
non-spontaneous f pent-2-yne
g 3-methylbut-1-yne
c E cell = +0.79V; h 3-methylpent-2-ene
ΔG = −460 kJ mol−1;
spontaneous
22 a anode: iodine cathode: hydrogen 5a H H CH3 H H H
b anode: oxygen cathode: hydrogen
c anode: chlorine cathode: hydrogen HCCCCCCH
d anode: oxygen cathode: hydrogen
e anode: oxygen cathode: silver HHHHHH
b H CH3 CH3 H H
HCCCCCH
23 a anode: 2H2O(l) → O2(g) + 4H+(aq) + 4e− H CH3 H H H
cathode: 2H2O(l) + 2e− → H2(g) + 2OH−(aq)
c H CH3 H CH3 H H
b anode: 2H2O(l) → O2(g) + 4H+(aq) + 4e−
cathode: Ag+(aq) + e− → Ag(s) HCCCCCCH
c anode: 2H2O(l) → O2(g) + 4H+(aq) + 4e− HHHHHH
cathode: 2H2O(l) + 2e− → H2(g) + 2OH−(aq)
d H CH3 H H
HCCCCH
CH3 H
568
6 a 3-methylpentane H OH H
b 3-methylpentane
c 2,2-dimethylbutane
7 a 1-bromobutane
b 2-bromo-2-methylbutane
c 2-chloropropane
d 1-iodo-3-methylbutane
8 HHHH
H C C C C OH HCCCH
HHH H H CH3 H
butan-1-ol H 2-methylpropan-2-ol
H OH H HHH
HCCCCH H C C C OH
HHHH H CH3 H
butan-2-ol* 2-methylpropan-1-ol
H HHH CH3
HCOCCC
H HOH
H HHH HCCCH
1-methoxypropane HHH
HH 2-methoxypropane
HH
H CCOCCH
HH HH
ethoxyethane
*Note: butan-2-ol also has optical isomers – see later in the Higher Level section, page 495.
9 a butanone
b 3-methylbutanal
c 2-methylpentanoic acid
d 3,3-dimethylbutanoic acid
10 Structure
a OHHH
CCCCH
H HHH O
HHHH
b
HCCCCC
H H H CH3 H
c HHOHH
HCCCCCH
HH HH
ANSWERS TO TEST YOURSELF QUESTIONS 569
11 a butyl propanoate 16 C5H12 + 8O2 → 5CO2 + 6H2O
b propyl methanoate
c propyl pentanoate 17 2C3H8 + 7O2 → 6CO + 8H2O
12 a 1-bromobutane 18 C2H6 + Cl2 → C2H5Cl + HCl
primary halogenoalkane 19 a 3 b 2 c 1
b 2-bromo-2-methylbutane tertiary halogenoalkane 20 Cl2 → 2Cl• initiation
Cl• + C2H6 → •CH2CH3 + HCl propagation
c 2-chloropropane secondary halogenoalkane •CH2CH3 + Cl2 → CH3CH2Cl + Cl• propagation
Cl• + Cl• → Cl2 termination
d 1-iodo-3-methylbutane primary halogenoalkane Cl• + •CH2CH3 → CH3CH2Cl termination
•CH2CH3 + •CH2CH3 → C4H10 termination
e 3,3-dimethylbutan-2-ol secondary alcohol
f 2-methylbutan-2-ol tertiary alcohol
g secondary amine
h primary amine
13
H HHH H CH3 H HH
CCCCH CCCH HCCCCH
HHHH
H HH H H
but-2-ene
but-1-ene 2-methylpropene
Note: cis–trans isomers are possible for but-2-ene – see page 490.
14
H HH H HH HH
HCCC CCH HCCCCCH HCCCCH
H HH H HH CH3 H
pent-1-yne pentm-2u-lytinplee bond is triple 3-methylbut-1-yne
15 a OH O
HH CCH HHH CH
HCCO H HCCCO
HH HH HHH
O O
H CCCH H CH
HCO HH H3C C O
CH3
H
b H OHHH O HHHH HH OHH
H C O C C C C HH C O C C C C HH C C O C C C H
H HHH HHHH HH HH
HO HHH O CH3 H H
O CH3 H
H C C O C C C HH C O C C C H HCOCCH
H HHH HHH CH3 H
O H CH3 H HO CH3 H H O CH3 H
HCOCCCH
HCCOCCH HCOCCCH
HHH H HH H HH
Note: one of these esters also exhibits optical isomerism.
570
21 a HHHH 27 a HH O
HCCCCH HCCC
HHHH HH OH
b H H H Br H H b HHHOH
HCCCCCH
HCCCCCCH
HHH H
HHHHHH c Not oxidised
c H H Cl Cl H d H CH3 H O
HCCCCCH HCCCC
HHHHH H CH3 H OH
d H OH H H 28 a O H H H c H H CH3 H
O
HCCCCH CCCCH HCCCCC
H H CH3 H
HHHH H H H CH3 H
22 a HH H H OH H H H b HHOHH
HH HCCCCCH
H C C C C C C H + H2O HCCCCCCH H CH3 HH
HHHHHH HHHHHH
b HH H + Br2 H H Br Br
CCCC HCCCCH
H H CH3 CH3 H
H CH3 CH3 H
23 a CH3 b CH3
H CH2
CC CH3 CH2
HH CC
Cl H
24 a H Cl
CC
Br H
b CH3
H CH2
CC
H CH3
25 CH3CH2CH2CH2OH + 6O2 → 4CO2 + 5H2O
26 a propan-1-ol primary
b pentan-2-ol secondary
c 2-methylbutan-2-ol tertiary
d 3,3-dimethylbutan-1-ol primary
ANSWERS TO TEST YOURSELF QUESTIONS 571
29 Alcohol Carboxylic acid Ester
aO
HHH CH
HCCCO
HHH
b H O H CH3 H
HCOCCCCH
H3C C CH3 H CH3 H
CH3
c H OH H OH
C C CH3
HCCCH
HHH HO H
d CH2OH O HHH
HH
HCCCH CCCCH
H CH3 H HO HHH
30 propyl methanoate
31 O
HHHH O
H C C C O+ conc. H2SO4 HHH CH
O
C H heat H C C C O +H H
HHH HO HHH
32 a SN2 34 a 2-chlorobutane
b 2-chloro-3-methylbutane
b SN1 c 2-chloro-2,4-dimethylpentane
c SN1 and SN2
d SN2 35 1-bromo-2-chlorobutane
33 a HHHHH 36 C6H5CH3 + HNO3 → CH3C6H4NO2 + H2O
HCCCCO
37 Formation of electrophile:
HNO3 + 2H2SO4 → NO2+ + H3O+ + 2HSO4−
HHHH NO2+ + NO2 NO2 + H+
H
bH electrophile
HOHH
HCCCCH 38 a propan-1-ol
b 2-methylpentan-1-ol
H CH3 H H c pentan-3-ol
d 3-methylbutan-1-ol
cH
HOH
HCCCH
HHH
d HHHHH
HCCCCO
H CH3 H H
572
39 a HH Br H OH H
H C C H Br2 H C C NaOH(aq) C C H
HH
UV heat
HH HH HH
bH CH3 H OH H HOH
C C + H2O conc. H2SO4 H C C C H Cr2O72–/H+ H C C C H
heat heat
HH HHH HH
c Cl H H OH H H O HH
H C C C H NaOH(aq) H C C C H Cr2O72–/H+ CCCH
heat heat under
HHH HHH reflux H O HH
d H Br H OH HO
HCCH HCCH HCC
H NaOH(aq) H Cr2O72–/H+ OH
heat
heat under
reflux
eH H H OH HO
C C + H2O conc. H2SO4 H C C H Cr2O72–/H+ H C C
heat heat
H OH
HH HH
conc. H2SO4 heat
O
H C CH3
H3C C O
H
40 Cl NaOH(aq) OH Cr2O72–/H+ OH
heat heat under O
A reflux
H OH conc. H2SO4
heat
CH
H
OH
B OCH
H
41 The alkene is 3-methylbut-1-ene: CH2=CHCH(CH3)2
c. H2SO4
CH2CHCH(CH3)2 + H2O ⎯⎯→ CH3CH(OH)CH(CH3)2
heat
Cr2O72−/ H+
CH3CH(OH)CH(CH3)2+ [O] ⎯⎯⎯→ CH3COCH(CH3)2 + H2O
heat / reflux
ANSWERS TO TEST YOURSELF QUESTIONS 573
42 H HH no
yes
H CCCCC H
no
H CH3 H H H yes HH HH
H H3C C C H H CCH
H C C C CH2CH3 C CH H C CH H
H H CH3 H CH3 H3C CH3
2,3-dimethylpent-2-ene H CH3 H
3,4-dimethylpent-2-ene HC CH H
HC CH3
HC CC H HC CH H
H CH3 H H CC H
1,2,3-trimethylcyclopropane yes H3C CH3 H CH3 H
1,3-dimethylcyclobutane yes C C C
H
HHC H H3C H C C
H3C H3C CH3
H H3C CH3 H H3C HH
C H
C C
C C HH C C C
HH HH CH3
H
43 a Z; b E; c Z; d E
44 a yes CH2OH CH2OH
CH HC
C2H5 OH HO C2H5
b no
c yes CH2OH CH2OH
CH HC
C2H5 CH3 H3C C2H5
d yes CH3 CH3
CH HC
(CH3)2HC Cl Cl CH(CH3)2
e no CHCHCH3 CHCHCH3
f yes
CH HC
H3C OH HO CH3
Note: chiral centres are shown in red.
45 Only H CH3
HC H
CC
H CC H
Cl H
HH
The others all have planes of symmetry.
46 a diastereomers; b enantiomers
574
Topic 11 c Non-linear relationship – as time increases the
1 Set 1 is more precise, because there is less variation temperature increases; as time increases the
from the mean.
increase in temperature per second gets larger
2 Experiment 2, because this value is closest to the
literature value. 11 a 1.0 cm3 s−1
b 17.5–18.4 gA−1
12 a 1 c2 e4
3 Measured b1 d1
value ± uncertainty Value you should quote 13 B, C, E and G
71.7 ± 0.2 71.7 ± 0.2 14 a 2840–3100 cm−1 by C–H
3.475 ± 0.01 3.48 ± 0.01 b 2840–3100 cm−1 by C–H; 3200–3600 cm−1 by
0.065 06 ± 0.001 0.065 ± 0.001
63.27 ± 5 63 ± 5 O–H
593.2 ± 30 590 ± 30 c 2840–3100 cm−1 by C–H; 1610–1680 cm−1 by
783.28 ± 100 800 ± 100
C=C
4 a3 c4 e4 d 2840–3100 cm−1 by C–H; 2400–3400 cm−1 by
b2 d3
O–H in COOH; 1700–1750 cm−1 by C=O;
5 a 6.79 (1000–1300 cm–1 by C–O)
b 0.000 079 8
c 0.005 00 15 a C2H4O
d 8.25 × 105 b C3H8O; C2H4O2
e 1.78 × 10−3 c C4H8O; C3H4O2
d C5H12O; C4H8O2; C3H4O3
6 0.345 ± 0.001 + 0.216 ± 0.002 = 0.561 ± 0.003
16 a CH3+
23.45 ± 0.03 − 15.23 ± 0.03 = 8.22 ± 0.06 b CO+; C2H4+
c C2H5+; HCO+
0.0034 ± 0.0003 + 0.0127 ± 0.0003 = 0.0161 ± 0.0006 d OCH3+; CH2OH+
e C3H7+;C2H3O+
1.103 ± 0.004 − 0.823 ± 0.001 = 0.280 ± 0.005 f COOH+; C2H5O+
g C3H7O+; CH2COOH+; HCOOCH2+
1.10 ± 0.05 + 17.20 ± 0.05 = 18.3 ± 0.1 h C6H5+
7 Value Percentage uncertainty (%) 17 a 3 peaks; 1 :2 :3
b 4 peaks; 3 :2 :2 :1
27.2 ± 0.2 0.74 c 3 peaks; 9 :2 :1
d 5 peaks; 1 :3 :1 :2 :3
0.576 ± 0.007 1.2 e 3 peaks; 3 :1 :6
f 5 peaks; 6 :1 :2 :2 :1
4.46 ± 0.01 0.22
7.63 × 10−5 ± 4 × 10−7 0.52 18 a two doublets
b one doublet; one triplet
8 a The absolute uncertainty is 0.05; the final value c one triplet; one quartet; one singlet
should be quoted as 9.22 ± 0.05. d one triplet; one quartet; one singlet
b The absolute uncertainty is 0.5; the final value 19 Using the values in Table 11.14:
should be quoted as 86.3 ± 0.5. a 0.5–5.0 ppm
b 2.0–2.5 ppm
9 4140 ± 40 J (assuming that the value for the specific c 3.7–4.8 ppm
heat capacity is exact)
10 a Linear relationship – as time increases the volume 20 Multiplicity of peak at δ = 4.0 ppm is 7
increases linearly; not a proportional relationship
H OH H
b Directly proportional – mass is directly
proportional to the current H CCCH
HHH
ANSWERS TO TEST YOURSELF QUESTIONS 575
Glossary
Terms in bold italic refer to keywords from online material
absolute scale of temperature Kelvin scale of temperature, ‘A’ is the frequency factor or pre-exponential factor
which starts at absolute zero; 1 °C is the same as 1 K and so and takes account of the frequency of collisions and the
0 °C is equivalent to 273 K orientation of the collisions
absolute zero the temperature at which everything would be atactic polymer a polymer that has side groups orientated
in its lowest energy state; 0 K or −273 °C randomly on both sides of the main chain
accuracy how close a measurement is to the actual value of a atom the smallest part of an element that can still be
particular quantity recognised as the element; in the simplest picture of the
acid deposition a more general term than acid rain; refers to atom, the electrons orbit around the central nucleus; the
any process in which acidic substances (particles, gases and nucleus is made up of protons and neutrons (except for a
precipitation) leave the atmosphere to be deposited on the hydrogen atom, which has no neutrons)
surface of the Earth; can be divided into wet deposition atom economy = molar mass of desired products × 100%
total molar mass of all reactants
(acid rain, fog and snow) and dry deposition (acidic gases
and particles) atomic number (Z) the number of protons in the nucleus of
acid dissociation constant (Ka) in general, for the an atom
dissociation of an acid HA: HA(aq) H+(aq) +A−(aq)
atomic radius half the internuclear distance between two
the expression for the acid dissociation constant is: atoms of the same element covalently bonded; atomic radius
[A−(aq)][H+(aq)]
Ka = [HA(aq)] is usually called ‘covalent radius’ in more advanced work; it is
also possible for an element to have a ‘van der Waals’ radius’
the higher the value of Ka, the stronger the acid Aufbau principle the process of putting electrons into atoms
activation energy (Ea) the minimum energy that colliding to generate the electronic configuration
species must have before collision results in a chemical average bond enthalpy the average amount of energy
reaction required to break one mole of covalent bonds, in a gaseous
actual yield the actual amount of product formed in the molecule under standard conditions;‘average’ refers to
reaction the fact that the bond enthalpy is different in different
addition polymerisation alkenes undergo addition molecules, and therefore the value quoted is the average
polymerisation, in which a large number of monomers are amount of energy to break a particular bond in a range of
joined together into a polymer chain; no other groups/ molecules; bond breaking requires energy (endothermic)
molecules are lost in the process ∆H +ve; bond making releases energy (exothermic) ∆H −ve
Avogadro’s constant (L) 6.02 × 1023 mol−1
addition reaction in organic chemistry, a reaction in which
a molecule is added to a compound containing a multiple Avogadro’s law equal volumes of ideal gases measured at the
bond without the loss of any other groups same temperature and pressure contain the same number of
alkali a base that is dissolved in water molecules
alkali metals the elements in group 1 of the periodic table base ionisation constant (Kb) consider the ionisation of a
weak base, B(aq) + H2O(l) BH+(aq) + OH−(aq):
allotropes different forms of the same element; e.g. diamond, [BH+(aq)][OH−(aq)]
[B(aq)]
graphite and fullerene are allotropes of carbon Kb =
alloy homogeneous mixture of two or more metals, or of a
metal with a non-metal the higher the value of Kb, the stronger the base
amphiprotic a substance that can donate a proton (acting as biochemical oxygen demand the amount of oxygen used
a Bronsted-Lowry acid) and accept a proton (acting as a by the aerobic microorganisms in water to decompose
Bronsted-Lowry base), e.g. HCO3–
the organic matter in the water over a fixed period of time
amphoteric a substance that can act as an acid and a base (usually 5 days) at a fixed temperature (usually 20 °C)
anabolism process of synthesising molecules needed by cells; biofuel a fuel produced from organic matter obtained from
requires energy plants, waste material etc.
analgesics drugs that reduce pain biomagnification the increase in concentration of a substance
anode the electrode at which oxidation occurs as it passes up a food chain
Arrhenius equation an equation showing the variation of bond enthalpy the enthalpy change when one mole of
the rate constant with temperature: covalent bonds, in a gaseous molecule, are broken under
k = Ae−Ea/RT
standard conditions
576
Born–Haber cycle an enthalpy level diagram breaking condensation polymers polymers formed when monomers,
down the formation of an ionic compound into a series of
simpler steps each containing two functional groups, join together
breeder reactor a nuclear reactor that produces more fissionable with the elimination of small molecules such as water or
material than it consumes
hydrogen chloride
Brønsted–Lowry acids and bases an acid is a proton (H+) conjugate acid–base pair these differ by one proton (H+);
donor; a base/alkali is a proton (H+) acceptor
when an acid donates a proton it forms the conjugate base
buffer solution one that resists changes in pH when small (CH3COO− is the conjugate base of CH3COOH); when
amounts of acid or alkali are added a base gains a proton it forms the conjugate acid (H3O+ is
the conjugate acid of H2O)
carbon footprint a measure of the total amount of greenhouse conjugated system a sequence of alternating single and double
gases (primarily carbon dioxide and methane) emitted as a
result of human activities bonds in a molecule - equivalent to a π delocalised system
catabolism breakdown of larger molecules into smaller ones continuous spectrum a spectrum consisting of all
with the release of energy
frequencies/wavelengths of light
catalyst a substance that increases the rate of a chemical
reaction without itself being used up in the reaction; a convergence limit the point in a line emission spectrum
catalyst acts by allowing the reaction to proceed by an
alternative pathway of lower activation energy where the lines merge to form a continuum; may be used
cathode the electrode at which reduction occurs to determine the ionisation energy
ceramics inorganic solid engineering materials that are neither
covalent bond the electrostatic attraction between a shared
metals nor polymers; they are usually described as inorganic
substances that contain at least one metallic and one pair of electrons and the nuclei of the atoms making up the
non-metallic element although substances such as silicon
carbide, diamond and graphite are also usually classified as bond
ceramics
chain reaction one initial event causes a large number of coordinate covalent bond a type of covalent bond in which
subsequent reactions – the reactive species is regenerated in
each cycle of reactions both electrons come from the same atom – also called a
chelate complex (chelate) a complex formed between a metal
ion and a polydentate ligand that results in the formation dative covalent bond
of a ring that includes the transition metal ion
chemical properties how a substance behaves in chemical coordination number the number of nearest neighbours for an
reactions
chiral centre a carbon atom with four different atoms or atom or ion in a crystal
groups attached to it; sometimes called an asymmetric
carbon atom cytoplasm the thick solution inside a cell but outside the
cis-trans isomerism two compounds have the same structural
formula, but the groups are arranged differently in space nucleus of the cell
around a double bond or a ring
closed system no exchange of matter with the surroundings degenerate describes orbitals with the same energy
collision theory a reaction can occur only when two
particles collide in the correct orientation and with E ≥ Ea delocalisation the sharing of a pair of electrons between
composite material mixture containing two or more different
materials present as distinct, separate phases; synthetic three or more atoms
composite materials consist of a reinforcing phase
embedded in a matrix diamagnetism caused by paired electrons – diamagnetic
concentration amount of solute dissolved in a unit volume of
solution; the volume that is usually taken is 1 dm3 substances are repelled slightly by a magnetic field
(one litre); the amount of solute may be expressed in g or
mol, so the units of concentration are g dm−3 or mol dm−3 diastereomers stereoisomers that are not mirror images of
condensation joining together of two molecules with the
formation of a covalent bond and the elimination of the each other
elements of water
dipole moment the product of one of the charges making
up a dipole and the distance between the charges; non-
polar molecules have a zero dipole moment
diprotic acid H2SO4 is a diprotic acid, as it can dissociate to
generate two protons per molecule:
H2SO4(aq) + H2O(l) → HSO4−(aq) + H3O+(aq)
HSO4−(aq) + H2O(l) SO42−(aq) + H3O+(aq)
drug any substance that, when applied to or introduced into
a living organism, brings about a change in biological
function through its chemical action
dynamic equilibrium macroscopic properties are constant;
rate of the forward reaction is equal to the rate of the
reverse reaction
ED50 the dose of a drug required to produce a therapeutic
effect in 50% of the test population (‘ED’ stands for
effective dose)
efficiency % efficiency = useful energy out × 100
total energy in
effusion the process by which a gas escapes through a very
small hole in a container
GLOSSARY 577
elastomers polymers that display rubber-like elasticity, they are the first electron affinity is exothermic for virtually all
elements
flexible and can be stretched to many times their original first ionisation energy the minimum amount of energy
required to remove an electron from a gaseous atom/the
dimensions by the application of a force; they will then energy required to remove one electron from each atom in
one mole of gaseous atoms under standard conditions
return to (nearly) their original size and shape once the fossil fuels fuels formed from things that were once alive and
have been buried underground for millions of years, e.g.
force is removed coal, oil and gas
free energy change (∆G) or Gibbs free energy change; ∆G is
electrolysis the breaking down of a substance (in molten related to the entropy change of the Universe and can be
defined using the equation:
state or solution) by the passage of electricity through it
∆G = ∆H − T∆S
electrolyte a solution, or a molten compound, that will
for a reaction to be spontaneous, ∆G for the reaction must
conduct electricity with decomposition at the electrodes be negative; ∆G is the standard free energy change
free radical a species (atom or groups of atoms) with an
as it does so; electrolytes contain ions that are free to move unpaired electron; free radicals are very reactive because of
this unpaired electron
towards the electrodes fuel cell a type of electrochemical cell that uses the reaction
between a fuel (such as hydrogen or methanol) and an
electronegativity a measure of the attraction of an atom in oxidising agent (e.g. oxygen) to produce electrical energy
directly; it uses a continuous supply of reactants from an
a molecule for the electron pair in the covalent bond of external source
functional group an atom or group of atoms that gives an
which it is a part organic molecule its characteristic chemical properties
genetic code how the four-base code in DNA determines the
electrophile a reagent (a positively charged ion or the sequence of 20 amino acids in proteins; each three base
codon codes for only one amino acid and this code is the
positive end of a dipole) that is attracted to regions of high same in all organisms – it is universal
genetically modified organisms (GMO) have genetic material that
electron density and accepts a pair of electrons to form a has been changed in some way by genetic engineering
giant structure bonding extends fairly uniformly throughout
covalent bond; an electrophile is a Lewis acid the whole structure; there are no individual molecules
green chemistry (also called ‘sustainable chemistry’) an approach
electroplating the process of coating an object with a thin to chemical research and chemical industrial processes that
seeks to minimise the production of hazardous substances
layer of a metal using electrolysis and their release to the environment
group vertical column in the periodic table
electrostatic attraction attraction between positive and Haber (Haber–Bosch) process an industrial process for the
manufacture of ammonia
negative charges half-life the time it takes for the number of radioactive nuclei
present in a sample at any given time to fall to half its value
element a substance containing just one type of atom halogens the elements in group 17 of the periodic table
heat a form of energy that flows from something at a higher
(although see isotopes) temperature to something at a lower temperature.
Hess’s law the enthalpy change accompanying a chemical
emission spectrum electromagnetic radiation given out reaction is independent of the pathway between the initial
and final states
when an electron falls from a higher energy level to a lower heterogeneous catalyst a catalyst that is in a different phase
(state) from the reactants
one; only certain frequencies of electromagnetic radiation heterogeneous mixture a mixture that does not have
uniform composition and consists of separate phases; can be
are emitted; each atom has a different emission spectrum separated by mechanical means
heterolytic fission a covalent bond breaks so that both
empirical formula the simplest whole number ratio of the electrons go to the same atom
elements present in a compound
enantiomers optical isomers
endothermic a chemical reaction in which heat is taken in
from the surroundings – the reaction vessel gets colder; ∆H
is positive for an endothermic reaction
end point of a titration the point at which an indicator
changes colour
energy density = energy released from fuel
volume of fuel consumed
enthalpy change (∆H) the heat energy exchanged with the
surroundings at constant pressure
entropy (S) a measure of how the available energy is
distributed among the particles; standard entropy (S ) is
the entropy of a substance at 100 kPa and 298 K; units are
J K−1 mol−1; ∆S is the entropy change under standard
conditions – a positive value indicates an increase in entropy,
i.e. the energy is more spread out (less concentrated)
equivalence point the point at which equivalent numbers of
moles of acid and alkali have been added in a titration
exothermic a chemical reaction that results in the release of
heat to the surroundings – the reaction vessel gets hotter;
∆H for an exothermic reaction is negative
first electron affinity enthalpy change when one electron is
added to each atom in one mole of gaseous atoms under
standard conditions:
X(g) + e− → X−(g)
578
homogeneous catalyst a catalyst that is in the same phase (state) ion a charged particle that is formed when an atom loses or
gains electron(s); a positive ion is formed when an atom
as the reactants loses (an) electron(s) and a negative ion is formed when an
atom gains (an) electron(s)
homogeneous mixture a mixture that has the same
ionic bonding the electrostatic attraction between oppositely
(uniform) composition throughout the mixture and charged ions
consists of only one phase ionic product constant (Kw) a modified equilibrium
constant for the dissociation of water:
homologous series a series of compounds with the same Kw = [H+(aq)][OH−(aq)]
Kw has a value of 1.0 × 10−14 at 25 °C
functional group, in which each member differs from the
isoelectronic describes species with the same number of
next by –CH2– electrons
homolytic fission a covalent bond breaks such that one
isotactic polymer a polymer in which all the side groups are on
electron goes back to each atom making up the original the same side of the polymer chain
covalent bond isotopes different atoms of the same element with different
mass numbers, i.e. different numbers of neutrons in the
Hund’s rule electrons fill orbitals of the same energy nucleus
(degenerate orbitals) so as to give the maximum number of kinetic energy the energy a body has because of its motion;
K.E. = 21mv2
electrons with the same spin
lattice structure regular 3D arrangement of particles
hybridisation the mixing of atomic orbitals when a Le Chatelier’s principle if a system at equilibrium is
compound forms to produce a new set of orbitals (the subjected to some change, the position of equilibrium will
shift in order to minimise the effect of the change
same number as originally), which are better arranged in Lewis acids and bases an acid is an electron pair acceptor; a
base is an electron pair donor
space for covalent bonding Lewis (electron dot) structure diagram showing all the
valence (outer shell) electrons in a molecule (or ion)
hydrocarbon compound containing carbon and hydrogen ligands negative ions or neutral molecules that use lone pairs
of electrons to bond to a transition metal ion to form a
only complex ion; coordinate covalent bonds (dative bonds) are
formed between the ligand and the transition metal ion
hydrogenation addition of hydrogen (H2) to a compound limiting reactant the reactant that is used up first in a
containing multiple bonds chemical reaction; when the number of moles of each
species is divided by their coefficient in the stoichiometric
hydrogen bonding an intermolecular force resulting from equation, the limiting reagent is the one with the lowest
number; all other reactants are in excess
the interaction of a lone pair on a very electronegative line spectrum the emission spectrum of an atom consists of
a series of lines that get closer together at higher frequency;
atom (N/O/F) in one molecule with an H atom attached only certain frequencies/wavelengths of light are present
lipid a broad class of biological molecules including steroids,
to N/O/F in another molecule; these forces may also triglycerides and phospholipids; mainly non-polar and
insoluble in water
occur intramolecularly liquid crystal a phase of matter in which the properties of a
compound may exhibit the characteristics of both a liquid
hydrolysis a reaction in which a covalent bond in a molecule and a solid
lone pair a pair of electrons in the outer shell of an atom
is broken by reaction with water; most commonly that is not involved in covalent bonding
London (dispersion) force intermolecular forces resulting
hydrolysis reactions occur when a molecule is reacted with from temporary (instantaneous) dipole–induced dipole
interactions
aqueous acid or aqueous alkali mass defect (∆m) is the difference between the mass of a
nucleus and the sum of the masses of the individual
ideal gas a theoretical model that approximates the behaviour nucleons
mass number (A) the number of protons + neutrons in the
of real gases; it can be defined in terms of macroscopic nucleus of an atom
properties (a gas that obeys the equation PV = nRT) or GLOSSARY 579
in terms of microscopic properties (the main assumptions
that define an ideal gas on a microscopic scale are that the
molecules are point masses – their volume is negligible
compared with the volume of the container – and that there
are no intermolecular forces except during a collision)
indicator an acid–base indicator has different colours
according to the pH of the solution; indicators are usually
weak acids (HIn); they dissociate according to the equation:
HIn(aq) H+(aq) + In−(aq)
Colour I Colour II
the ionised (In−) and un-ionised (HIn) forms must have
different colours. Indicators may also be weak bases
initiation step a step that starts off a chain reaction; it
involves an increase in the number of free radicals
intermolecular forces forces between molecules
internal energy the total amount of energy (kinetic and
potential) in a sample of a substance
intramolecular forces forces within a molecule – usually
covalent bonding
iodine number a measure of the degree of unsaturation in a fat
or oil; it is the number of grams of iodine that reacts with
100 g of fat or oil
Markovnikov’s rule when H–X adds across the double bond nucleophile a molecule/negatively charged ion, possessing
of an alkene, the H atom becomes attached to the C atom
that has the larger number of H atoms already attached a lone pair of electrons, which is attracted to a more
matter something that has mass and occupies space positively charged region in a molecule (region with lower
Maxwell–Boltzmann distribution a graph showing the
electron density) and donates a lone pair of electrons to
distribution of molecular kinetic energies in a sample of gas
at a particular temperature form a covalent bond; a nucleophile is a Lewis base
medicine something that treats, prevents or alleviates the
symptoms of disease nucleophilic substitution a nucleophile replaces an atom/
metabolism chemical reactions that go on in cells – it involves
the breakdown of molecules with the release of energy and group in a molecule
the synthesis of molecules that are required by cells
metallic bonding the electrostatic attraction between the octane number a measure of the tendency of a fuel to not
positive ions in a metallic lattice and the delocalised electrons
Michaelis constant (Km) in enzyme kinetics is the undergo auto-ignition (not cause knocking) in an engine;
concentration of substrate when the rate is equal to one
half of Vmax (maximum rate) the higher the octane number, the lower the tendency to
molar volume the volume occupied by one mole of a gas;
the molar volume of an ideal gas at STP is 22.7 dm3 mol−1 undergo autoignition/cause knocking
mole the amount of substance that contains the same
number of particles (atoms, ions, molecules, etc.) as there opiates natural narcotic (sleep-inducing) analgesics derived
are carbon atoms in 12 g of carbon-12 (6.02 × 1023)
molecular formula the total number of atoms of each from the opium poppy
element present in a molecule of the compound; the
molecular formula is a multiple of the empirical formula optical isomersism when molecules have the same
molecular self-assembly a bottom-up approach to producing
nanoparticles, where molecules come together reversibly molecular and structural formula, but groups are arranged
and spontaneously to create a larger structure
molecularity the number of ‘molecules’ that react in a differently in space and the individual optical isomers are
particular step (usually the rate-determining step) in a
chemical reaction non-superimposable mirror images of each other; optical
molecule an electrically neutral particle consisting of two or
more atoms chemically bonded together isomers rotate the plane of plane-polarised light in opposite
monomer a molecule from which a polymer chain may be
built up, e.g. ethene is the monomer for polyethene directions
monoprotic acid HCl is a monoprotic acid as it dissociates
to form one proton per molecule orbital a region of space in which there is a high probability
nanotechnology the production and application of structures,
devices and systems at the nanometre scale; generally, of finding an electron; it represents a discrete energy level;
involves man-made particles or structures that have at least
one dimension smaller than 100 nm there are s, p, d and f orbitals; any orbital can contain a
nematic liquid crystal phase molecules point, on average, in the
same direction but are positioned randomly relative to each maximum of two electrons
other (no positional order)
noble gases the elements in group 18 of the periodic table; order of a reaction the power of the concentration of a
also sometimes called the ‘inert gases’
non-renewable energy sources sources of energy that are finite – particular reactant in the experimentally determined rate
they will eventually run out, e.g. coal
nuclear binding energy (∆E) is the energy required to break equation
apart a nucleus into individual protons and neutrons
nuclear fission the breakdown of a larger nucleus into two osmosis movement of water (or other solvents) across a
smaller fragments of comparable masses
nuclear fusion the joining together of smaller nuclei to make semipermeable membrane from a less concentrated
a larger one
solution to a more concentrated one
580
overall order of reaction the sum of the powers of the
concentration terms in the experimentally determined rate
equation
oxidation number a purely formal concept that regards all
compounds as ionic and assigns charges to the components
accordingly; it provides a guide to the distribution of
electrons in covalent compounds
oxidation loss of electrons or increase in oxidation number
oxidising agent (oxidant) oxidises other species and, in the
process, is itself reduced; an oxidising agent takes electrons
away from something
paramagnetism caused by unpaired electrons – paramagnetic
substances are attracted by a magnetic field
Pauli exclusion principle two electrons in the same orbital
must have opposite spins
percentage error =|experim|enactaclevpatleude−vaalcucee|pted value|×100
percentage uncertainty = uncertainty × 100
measured value
when multiplying or dividing quantities with uncertainties,
the percentage uncertainties should be added to give the
percentage uncertainty of the final value
period horizontal row in the periodic table
pH a measure of the concentration of H+ ions in an aqueous
solution; it can be defined as the negative logarithm to base
10 of the hydrogen ion concentration in aqueous solution:
pH = −log10 [H+(aq)]
pH meter an electronic device for measuring the pH of a rate-determining step the slowest step in a reaction
solution
mechanism
pH range of an indicator the pH range over which
intermediate colours for an indicator can be seen rate constant (k) a constant of proportionality relating the
physical properties properties such as melting point, concentrations in the experimentally determined rate
solubility and electrical conductivity, relating to the
physical state of a substance and the physical changes it can expression to the rate of a chemical reaction; the rate
undergo
constant is only a constant for a particular reaction at a
pi (π) bond bond formed by the sideways overlap of parallel
p orbitals; the electron density in the pi bond lies above and particular temperature
below the internuclear axis
rate expression or rate equation an experimentally
plane-polarised light light that vibrates in one plane only;
optical isomers rotate the plane of plane-polarised light in determined equation that relates the rate of reaction to the
opposite directions
concentrations of substances in the reaction mixture, e.g.:
plasma a fully or partially ionised gas consisting of positive rate = k[A]m[B]n
ions and electrons
rate of reaction the speed at which reactants are used up
plasticisers small molecules that are added to a polymer to
increase its flexibility or products are formed; or, more precisely, the change in
polar molecule molecule in which one end is slightly concentration of reactants or products per unit time
positive relative to the other; whether a molecule is polar
or not depends on the differences in electronegativity of redox reaction a reaction that involves both oxidation and
the atoms and the shape of the molecule
reduction; if something is oxidised, something else must be
polydentate ligand a ligand that binds to a transition metal ion
through more than one donor atom reduced
polymers long-chain molecules, usually based on carbon, reducing agent (reductant) reduces other species and, in the
which are formed when smaller molecules (monomers)
join together process, is itself oxidised; a reducing agent gives electrons to
precision relates to the reproducibility of results; if a series of something
readings is taken with high precision, it indicates that the
repeated values are all very close together and close to the reduction gain of electrons or decrease in oxidation number
mean (average) value
relative atomic mass (Ar) the average of the masses of the
primary cell a cell (battery) that cannot usually be recharged isotopes in a naturally occurring sample of the element
using mains electricity – the reaction in the cell is non- 1
reversible relative to the mass of 12 of an atom of carbon-12
primary structure of a protein the linear sequence of amino relative formula mass if a compound contains ions, the
acids in a polypeptide chain
relative formula mass is the mass of the formula unit
propagation step a step in a free radical substitution reaction 1
that involves production of products and no change in the relative to the mass of 12 of an atom of carbon-12
number of free radicals
relative molecular mass (Mr) the mass of a molecule of
prosthetic group a non-peptide part of a protein that is bound 1
tightly to the protein and is required for correct function; a compound relative to the mass of 12 of an atom of
for example, the heme group in haemoglobin
carbon-12; the Mr is the sum of the relative atomic masses
racemic mixture an equimolar mixture of the two for the individual atoms making up a molecule
enantiomers of a chiral compound; it has no effect on
plane-polarised light renewable energy sources sources of energy that are naturally
radioisotopes radioactive isotopes replenished – they will not run out, e.g. solar energy, wind
random uncertainty uncertainty in a measurement
power
due to the limitations of the measuring apparatus and
other uncontrollable variables that are inevitable in any repeating unit (repeat unit) of a polymer the basic unit
experiment; the effects of random uncertainties should
mean that the measurements taken will be distributed from which the whole polymer chain can be made up
either side of the mean, i.e. fluctuations will be in both
directions; the effect of random uncertainties can be resonance hybrid the actual structure of a molecule/ion for
reduced by repeating the measurements more often, but
random uncertainties can never be completely eliminated which resonance structures can be drawn can be described
as resonance hybrid made up of contributions (not
necessarily equal) from all possible resonance structures
resonance structure one of several Lewis structures that can
be drawn for some molecules/ion
salt bridge completes the circuit in a voltaic cell by providing
an electrical connection between two half cells, allowing
ions to flow into or out of the half cells to balance out
the charges in the half cells; the salt bridge contains a
concentrated solution of an ionic salt such as KCl
saturated compounds organic compounds containing only
single bonds
second electron affinity the enthalpy change for the process:
X−(g) + e− → X2−(g)
the second electron affinity is always endothermic
second ionisation energy the enthalpy change for the
process:
M+(g) → M2+(g) + e−
GLOSSARY 581
secondary cell a cell (battery) that can be recharged using standard enthalpy change of formation (∆Hf ) the enthalpy
change when one mole of the substance is formed from its
mains electricity and is often called a rechargeable battery;
the chemical reactions in a rechargeable battery are elements in their standard states under standard conditions;
reversible and can be reversed by connecting them to an ∆H f for any element in its standard state is zero
standard enthalpy change of hydration (∆H hyd) the enthalpy
electricity supply
change when one mole of gaseous ions are surrounded
side effect an unintended secondary effect of a drug on the
body; it is usually an undesirable effect by water molecules to form an ‘infinitely dilute solution’
sigma (σ) bond bond formed by the axial (head-on) overlap under standard conditions:
of atomic orbitals; the electron distribution in a sigma bond Na+(g) ∆Hhyd Na+(aq)
⎯⎯⎯⎯→
lies mostly along the axis joining the two nuclei excess H2O
SN1 a unimolecular nucleophilic substitution reaction – only an exothermic process
one species involved in the rate-determining step
standard enthalpy change of neutralisation (∆Hn) the
SN2 a bimolecular nucleophilic substitution reaction – two enthalpy change when one mole of H2O molecules are
species involved in the rate-determining step formed when an acid (H+) reacts with an alkali (OH−)
solute a substance that is dissolved in another (the solvent) under standard conditions, i.e.:
H+(aq) + OH−(aq) → H2O(l)
solution that which is formed when a solute dissolves in a
standard enthalpy change of reaction (∆Hr ) the enthalpy
solvent change (heat given out or taken in) when molar amounts
solvent a substance that dissolves another substance (the
solute); the solvent should be present in excess of the solute of reactants as shown in the stoichiometric equation react
specific energy = energy released from fuel together under standard conditions to give products
mass of fuel consumed
the enthalpy change of neutralisation is always exothermic
specific heat capacity the energy required to raise the
standard enthalpy change of solution (∆Hsol) the enthalpy
temperature of 1 g of substance by 1 K (1 °C) – units change when one mole of solute is dissolved in excess
J g−1 °C−1; it can also be defined in terms of ‘unit mass’ with
solvent to form a solution of ‘infinite dilution’ under
different units
standard conditions, e.g.:
spectrochemical series a series of ligands arranged in order
NH4NO3(s) ⎯exc⎯es⎯s H⎯2→O NH4+(aq) + NO3−(aq)
of the extent to which they cause splitting of d orbitals in a ‘infinite dilution’ means that any further dilution of the
solution produces no further enthalpy change, i.e. the
transition metal complex ion
spontaneous reaction one that occurs without any outside
influence, i.e. no input of energy; ∆G is negative solute particles are assumed not to interact with each other
stability usually refers to the relative energies of reactants in the solution; the enthalpy change of solution may be
and products – if the products are at lower enthalpy exothermic or endothermic
(energy) than the reactants, then they are more stable; it is standard enthalpy change of vaporisation (∆Hvap) the
energy needed to convert one mole of a liquid to vapour
also possible to define kinetic stability
under standard conditions
standard cell potential a standard cell potential is the EMF
standard hydrogen electrode the standard half-cell relative to
(voltage) produced when two half-cells are connected
which standard electrode potentials are measured
under standard conditions.This drives the movement of
electrons through the external circuit from the negative standard lattice enthalpy (∆Hlatt) the enthalpy change when
one mole of ionic compound is broken apart into its
electrode to the positive electrode
standard conditions a common set of conditions used to constituent gaseous ions under standard conditions, e.g. for
compare enthalpy changes; the pressure is 100 kPa and for NaCl:
reactions involving solutions, all solutions should have a NaCl(s) → Na+(g) + Cl−(g) ∆Hlatt = +771 kJ mol−1
concentration of 1 mol dm–3; a temperature is not stated
lattice enthalpy can be defined in either direction, i.e. as the
in the definition and should be specified – when it is not
making or breaking of the lattice
specifed we will assume in this course that it is 298K (25 °C)
standard state the pure substance at 100 kPa and a specified
standard electrode potential the emf (voltage) of a half-cell temperature (assume 298 K if one is not given). It is
connected to a standard hydrogen electrode, measured often used to refer the state in which a substance exists
under standard conditions; all solutions must be of under standard conditions, e.g. for iodine it is I2(s) but for
concentration 1 mol dm−3 nitrogen it is N2(g)
standard temperature and pressure (STP) 273 K, 100 kPa
standard enthalpy change of atomisation (∆Hat ) this is state symbols used to indicate the physical state of an
the enthalpy change when one mole of gaseous atoms is element or compound; these may be either written
formed from an element under standard conditions
standard enthalpy change of combustion (∆Hc ) the as subscripts after the chemical formula or in normal
enthalpy change (heat given out) when one mole of a
type (aq) = aqueous (dissolved in water); (g) = gas;
substance is completely burnt in oxygen under standard
(l) = liquid; (s) = solid
conditions
582
stereoisomers molecules with the same molecular formula thermosetting polymer a prepolymer in a soft solid or viscous
state that changes irreversibly into a polymer network
and structural formula but the atoms are arranged (thermoset) by curing
differently in space; cis-trans isomers, conformational titration a technique that involves adding measured amounts
of a solution (from a burette) to another solution to
isomers and optical isomers are stereoisomers determine the amounts that react exactly with each other
strong acid an acid such as HCl, H2SO4, HNO3 that tolerance when the body becomes less responsive to the
dissociates completely in aqueous solution: effects of a drug – larger and larger doses are needed to
HCl(aq) → H+(aq) + Cl−(aq) produce the same effect which means that the patient may
be at higher risk of toxic side effects
strong acids are also strong electrolytes
transition elements the elements in the central part of
strong base a base that ionises completely in aqueous the periodic table; there are various ways of defining a
transition element; the definition used in the IB course
solution; strong bases are the group 1 hydroxides (LiOH, is ‘an element which forms at least one stable ion with a
partially filled d subshell’
NaOH, etc.) and Ba(OH)2; strong bases are also strong
electrolytes transition state (activated complex) the highest energy
species on the reaction pathway between reactants and
strong electrolyte a substance that dissolves in water with products; the highest point on potential energy profile
complete ionisation; e.g. NaCl separates completely into unit cell the simplest repeating unit from which a whole
crystal can be built up
its ions when it dissolves in water; strong acids are strong
unsaturated compounds organic compounds containing
electrolytes because they dissociate completely into ions multiple bonds; this term is often just applied to
compounds containing C=C and C=−C, but it is more
structural isomers two or more compounds that have the widely applicable and compounds containing e.g. C=O are
also unsaturated
same molecular formula but different structural formulas,
valence shell electron pair repulsion (VSEPR) theory a
i.e. the atoms are joined together in a different way technique for working out the shapes of molecules/ions; pairs
of electrons in the valence (outer) shell of an atom repel each
substitution reaction a reaction in which one atom or group other and therefore take up positions in space to minimise
these repulsions, i.e. to be as far apart in space as possible
is replaced by another atom or group
van der Waals’ forces the collective name given to the forces
superconductors materials that have zero electrical resistance between molecules and includes London (dispersion)
forces, dipole−dipole interactions and dipole−induced
below a critical temperature dipole interactions but not hydrogen bonding and ion-
dipole interactions
sub-energy level a group of degenerate orbitals in an atom; a
vapour pressure the pressure exerted by a vapour in
p subshell is made up of the px, py and pz orbitals equilibrium with a liquid (or a solid)
syngas (synthesis gas) a mixture of carbon monoxide and
viscosity a measure of the resistance of a liquid to flow
hydrogen volatility how readily a substance evaporates
water of crystallisation water that is present in definite
systematic error an error introduced into an experiment
proportions in the crystals of hydrated salts, e.g.
due to the apparatus used or the procedure; systematic CuSO4·5H2O.The water may or may not be directly
bonded to the metal (in hydrated copper sulfate four water
errors result in a loss of accuracy, i.e. the measured value molecules are bonded to the copper ion and one is not)
weak acid an acid such as a carboxylic acid (ethanoic
being further away from the true value; systematic errors acid, propanoic acid, etc.) or carbonic acid (H2CO3) that
dissociates partially in aqueous solution:
are always in the same direction; the effect of a systematic
CH3COOH(aq) H+(aq) + CH3COO−(aq)
error cannot be reduced by repeating the readings weak acids are also weak electrolytes
weak base a base that ionises partially in aqueous solution,
TD50 the dose of a drug required to produce a toxic effect in e.g. ammonia and amines:
NH3(aq) + H2O(l) NH4+(aq) + OH−(aq)
50% of the test population (‘TD’ stands for toxic dose) weak bases are also weak electrolytes
temperature a measure of the average kinetic energy of
particles
termination step the step that ends a chain reaction –
involves a decrease in the number of free radicals
theoretical yield the maximum possible amount of product
formed
therapeutic effect a desirable and beneficial effect of a drug;
one that alleviates symptoms or treats a particular disease
therapeutic index (TI) the ratio of the toxic dose to the
therapeutic dose of a drug
TI =ELDD5500 or TI = TD50
ED50
therapeutic window the range of dosage of a drug between the
minimum required to cause a therapeutic effect and the
level which produces unacceptable toxic effects
thermoplastic a type of polymer that softens when it is heated
and hardens when it is cooled; it can be repeatedly heated
and cooled and remoulded into different shapes
GLOSSARY 583
weak electrolyte a substance that is partially ionised when
dissolved in water
xenobiotics compounds that are present in living organisms
but should not normally be found there
yield the amount of product obtained from a chemical
reaction; see also actual yield and theoretical yield
zeolites usually refers to an aluminosilicate, porous crystalline
mineral (either natural or synthetic)
584
Index
absolute temperature scale 29 oxidation of 458–60 ammonium chloride 355
absolute uncertainty 516 primary 439 ammonium ethanoate 356
absolute zero 29 reactions of 458–62 ammonium ion 131, 140
accuracy 510, 512 secondary 439 amphiprotic substances 309
acid deposition 325–7 solubility of 158 amphoteric oxides 101
tertiary 439 amphoteric substances 309
methods of dealing with 326 see also ethanol analogue instruments 509
problems associated with 326 aldehydes 458, 459 anodes 387, 391
see also acid rain naming 436 aqueous solutions, electrolysis of 406–15
acid dissociation constant (Ka) 328–30 reduction of 481 aromatic compounds 426, 487
calculation of 333 aliphatic compounds 426 Arrhenius equation 268–71
relationship between Kb and 339 alkali metals/group 1 metals 97–8 asymmetric carbon atom 495
see also pKa atomic and ionic radii for 90 atomic emission spectra 65
acid rain 103, 327 melting points 97 atomic number (Z) 57
see also acid deposition reactions of 98 atomic orbitals 167
acidic oxides 101, 102 atomic radii 88–9
acidified water 407, 412 with oxygen 98 atomisation, enthalpy change of 207
acids with water 98 atoms 3, 56–7
acid−base titrations 341–53 alkalis 313
acidity due to positive ions in solution see also bases electrons in 62–6, 72–9
356–7 alkanes 422–3, 424, 447–51 energy levels in 63–6
Brønsted−Lowry 308–9 boiling points 149–50 see also energy levels; orbitals; shells,
buffer solutions 358–60 combustion of 447 electron
concentration 324 distinguishing between alkenes and 454 Aufbau principle 67–8, 70–1
conduction of electricity 323 naming 428–30 average bond enthalpy 208
diprotic 320 reactions of 447 Avogadro’s constant 9
monoprotic 320 with halogens 447–50 Avogadro’s law 27
pH 323 unreactivity of 447
properties of 312–4 see also methane (CH4) back titrations 47–8
alkenes 426 balancing equations 4–7, 375–7
reactions with bases and alkalis 313 addition polymerisation 455–7 base ionisation (dissociation) constant (Kb)
reactions with carbonates and addition reactions 451–4
hydrogencarbonates 312 distinguishing between alkanes and 454 330–1
reactions with metals 312 electrophilic addition reactions of 473–7 relationship between Ka and 339
rate of reaction 323 hydrogenation of 453 see also pKb
relationship between the strength of an naming 430–1 bases 321–2
acid and the strength of its conjugate base reactions of 451–4 Brønsted−Lowry 308–9
339–40 with halogens 452 buffer solutions 360
strength 324 with hydrogen 452 properties of 312–4
strong see strong acids with hydrogen halides 452 reactions with acids 313
weak see weak acids with water 452–3 see also base ionisation (dissociation)
see also acid dissociation constant (Ka) see also ethene constant (Kb)
actinoide elements 85 alkynes 431, 451 basic oxides 101
activated complex 466 allotropes of carbon 144–7 BeCl2 130
activation energy 246, 268–71 alloys 161 bent molecule shape 139, 140, 141
in the Arrhenius equation 268–71 aluminium 101, 393 benzene 172–3, 426, 431, 487
and frequency factor values 271 aluminium fluoride 122 carbon−carbon bond lengths 445
activity series 380–1, 404 aluminium oxide 101, 392, 412 delocalised π system 478
actual yield, of chemical reactions 23 amines 426 index of hydrogen deficiency (IHD) of
addition polymerisation 455–7 primary 440 525
addition reactions 451 secondary 440 nitration of 479–80
adducts 131 tertiary 440 reactions of 446, 454
air-conditioning 2 ammonia (NH3) 140, 179, 322 structure of 133, 444–6, 478
alcohols 424 hydrogen bonding in 153 beryllium 78, 79, 130
boiling points of 427 as a Lewis base 310, 311 BF3 (boron(III) fluoride) 130, 143, 179
combustion of 458 production of 109, 232, 302 biochemical/biological oxygen demand
naming 434–5 shape of 138
(BOD) 382–3
INDEX 585
boiling points 149–50, 154–6 carbocations 468, 474, 475, 476 ClF3 (chlorine(III) fluoride) 166
of alcohols 427 primary 470, 474, 475 closed systems 278, 280
of ethers 427 secondary 470, 474, 475 CO32– (carbonate) ions 133, 171, 172
of group 16 hydrides 152–3 in SN1 mechanism 469 collision theory 246–51
and homologous series 427 stability of 470 colorimeters 245
of hydrocarbons 149–50 tertiary 470, 475 colour wheel 111
of ionic compounds 123 coloured complexes 110–4
of a pair of isomers 492 carbon 125, 176 combustion
relative molecular/atomic mass and 149 allotropes of 144–7
see also buckminsterfullerene; diamond; of alkanes 447
bomb calorimeters 191 graphite complete 447
bond angles 140 enthalpy changes of 189–91
bond enthalpies/bond energy 207–14 carbon dioxide (CO2) 139, 143, 179 incomplete 447
covalent bonding in 127 standard enthalpy change 188
calculation from enthalpy change of Lewis structures for 136 complex ions 105, 107, 113
reaction 213 in rate of reaction experiments 241–3 charge 109
enthalpy changes for reactions calculated oxidation number 109, 370
from 209–10 carbon monoxide 131–2 transition metal ion 108–9, 110, 111, 311
using cycles in calculations involving carbonates, reactions of acids with 312, 323 see also ligands
210–2 carbonic acid 320, 321 complex multiplets 543
bonding 119 carbonyl compounds 436 compounds 3
hybridisation 176–9 percentage composition of 12–4
hydrogen 152–4, 157 see also aldehydes; ketones concentration 287
metallic 160–1 carboxylic acids 437, 443, 458, 459 and equilibrium 285, 292, 297
see also covalent bonding; intermolecular of ions 42
forces; ionic bonding naming 436 of solutions 40
Born−Haber cycles 216–20 reduction of 482 of very dilute solutions 41–2
boron 78, 79 see also ethanoic acid condensation 281, 463
boron(III) fluoride (BF3) 130, 143, 179 catalysts 250, 264–6 conformational isomers 494–5
Boyle’s law 32–3 effect on the value of the rate constant conjugate acid−base pairs 308
BrF5 (bromine(V) fluoride) 166 271 conjugate acids 308
brittleness 161 and equilibrium 292 conjugate bases 308
bromine 148, 151 and the equilibrium constant 290–1 conservation of mass 18–9
reaction with ethane 450 homogeneous 265 continuous data 522
reaction with ethene 476–7 catalytic hydration 453 continuous spectra 63
bromine water test 454 catenation 422 convergence 66
bromine(V) fluoride (BrF5) 166 cathodes 388, 391 convergence limit 72
3-bromo-2-chloropentane 433 cell potentials 393–4, 397 coordinate covalent bonds 131–2, 162, 167,
2-bromo-5-methylhexane 433 CFCs (chlorofluorocarbons) 2, 174–5
bromonium ion 477 chain reactions 449 310
Brønsted−Lowry acids and bases 308, 310, changes of state 1, 2 copper 71, 410
charge copper(II) sulfate 386
311 on ions 109, 122
buckminsterfullerene 146–7 mass:charge ratio 61, 530 anode product in the electrolyis of
buffer solutions 358–61 on sub-atomic particles 56 408–9
butane 423 see also nuclear charge electrolysis using copper electrodes
butanoic acid 527 charge density 113 410–1
butanone 534 Charles’ law 33 electrolysis using inert electrodes 409–10
butenedioic acid 154 chemical equations 3–4 hydrated 45
chemical equivalence 540 covalent bonding 125–30, 176–7
C6H4Cl2, isomers of 446 chemical properties 3 multiple 126–7
Cahn–Ingold–Prelog priority rules (CIP) chemical reactions 4 single 125–6
calculating the yield of 23–4 covalent compounds 119, 130–79
491 chemical shift 534, 536–8, 543 intermolecular forces 148–59
calcium carbonate 22, 278, 284 chiral centre 495 physical properties of 159
calcium oxide 278 chiral molecules 495 shapes of 137–47
calculations chlorine 89, 96 Cr2O72– 370
chlorine(III) fluoride (ClF3) 166 current 413
involving acids and bases 328–40 chloroethane 532, 539 cyclic compounds 492
involving moles and masses 18–26 chlorofluorocarbons (CFCs) 2, 174–5 optical isomerism and 496–7
involving solutions 39–51 chloromethane 448 cycloalkanes 492–3
involving volumes of gases 27–39 chromium 71 cyclohexane 494
uncertainties in 515–9 cis-trans isomerism 489–90 cyclohexene 445
yield of a chemical reaction 23–4 properties of 492
calorimetry 189 in substituted cycloalkanes 492–3
see also geometrical isomers
586
d-block elements 104–10 electrolytes 323, 391 equilibrium concentrations 286
see also complex ions; transition metals electrolytic cells 390–2 equilibrium constants (Kc) 286–92
electromagnetic radiation 63, 72–3
decimal places 514, 515 electromagnetic spectrum 62–3 calculation of 292–9
degree of unsaturation see index of electron affinity 93–5, 215 catalysts and 290–1, 292
electron configurations 67–71 concentration and 292, 297–8
hydrogen deficiency (IHD) effect of temperature on 289–90, 292
delocalisation 170–3, 444 periodic table and 87–8 and Gibbs free energy 300–1
electron domains 137 pressure and 292
in metals 160 electron dot structures see Lewis structures and rate of reaction 291
see also pi (π) delocalised systems electron-releasing effect 470, 475 values of 288–9
dependent variables 522 electronegativity 95–6, 128–30 see also acid dissociation constant (Ka);
desalination 7 electrons 56 base ionisation (dissociation) constant
diamagnetism 107 (Kb); ionic product constant (Kw)
diamond 145 in atoms 62–6, 72–9 equilibrium law 286, 292–301
diastereomers 499 electron−electron repulsion 94 equivalence point 342, 349, 351
digital instruments 509 lone pairs of 126, 140 strong acid−weak base 347
dimethylamine 330 valence 87 weak acid−strong base 345
2,2-dimethylbutane 429 see also energy levels; orbitals; shells, weak acid−weak base 348
dipole−dipole interactions 150, 155 electron errors 510, 511–2
dipole moment 143 electrophiles 473, 479 percentage 511
discrete data 522 electrophilic addition reactions 473–7 systematic 510, 511–2
disorder 225, 227 electrophilic substitution reactions 478–80 see also uncertainties
dispersion forces see London (dispersion) electroplating 411 esterification 462–3
electrostatic attraction 122, 123 esters 438–9, 443, 462–3
forces elements 3, 4 naming 436, 438
displacement reactions 100, 381 emission spectra 63, 64, 65 uses of 438
dissociation empirical formulas 14–7, 422 ethanal 150–1
enantiomers 496, 498, 499 ethane, reaction with bromine 450
of acids 319, 321, 328 see also optical isomerism ethanoic acid 320, 328, 443
of water 316 end point 349, 351, 352 ethanol 157, 439, 453
dissolved oxygen 382–3 endothermic reactions 187, 211, 225, 290 combustion of 189–90
double bonds 126, 127, 170 definition of 185 hydrogen bonding in 154
in ethene 169 enthalpy change of 186 ethene
and shapes of molecules 137 potential energy diagrams (profiles) 213 bonding in 127, 169, 177
ductility 161 energy reaction with bromine 476–7
dynamic equilibrium 279 activation see activation energy reaction with hydrogen bromide 473–4
for bond breaking 207 ethers 154
E/Z isomers 490–1, 492, 493 for bond making 207 boiling points of 427
effective nuclear charge 75, 78 ionisation see ionisation energy naming 435
electrical conductivity, of ionic compounds see also enthalpy changes (∆H); Gibbs free 4-ethyl-2-methylhexane 430
energy; heat ethyl methanoate 443
124 energy cycles 215–24 ethylamine 322
electricity, conduction in electrolytic cells energy levels 62, 63–6 ethylbenzene 487
enthalpy changes (∆H) 186, 187 ethyne 127
391 of atomisation 207 bonding in 178
electrochemical cells 386–93 bond enthalpies 207–14 shape of 142
of combustion 189–91 evaporation 280–1
electrolytic cells 390–3 of formation 203–6 exothermic reactions 187, 211, 232,
standard electrode potentials 393–406 and Hess’s law 196–202
voltaic cells 386–90 of hydration 223, 224 289–90
electrochemical series 404 of neutralisation 191 activation energy 246
electrode potentials 404 of reactions 205, 209–10 definition of 185
electrodes 387–8, 391 of solutions 191–5, 223 enthalpy change of 186
and nature of products formed during enthalpy level diagrams 201 potential energy diagrams (profiles) 213
electrolysis 409–11 entropy 225
negative 391 of the Universe 228–9 first electron affinity 215
polarity of 398–9 entropy change 226–7 first ionisation energy 73, 92–3, 215
positive 391 equilibrium 278–92 first law of thermodynamics 228
electrolysis characteristics of 280 first-order reactions 258–9
of aqueous solutions 406–15 physical 280–1 fluorine 96, 128, 149
effect of concentration of the electrolyte position of 281–5 formal charge 162–3
on the products at the electrodes 409 and rate of reaction 279
formation of halogens during 408 INDEX 587
of molten salts 390–2
quantitative 413–5
and standard free energy change 412
of water 412
formation, standard enthalpy change of H3O+ (hydroxonium ion) 141 hydrogen peroxide 370
203–6 Haber process 109, 232, 302 hydrogencarbonate ion 321
half-equations 368, 374–8 hydrogencarbonates, reactions of acids with
formulas
empirical 14–7 balancing half-equations 312
of ionic compounds 122 in acidic solution 376–7 hydrolysis 465
of ions 121–2 in neutral solution 375
molecular 14–7 I3– (triiodide ion) 166
for solving moles questions involving for redox equations 377 ideal gas equation 36–8
masses 21 half-life 259 ideal gases 27, 32–5
for solving moles questions involving halide ions 100
volumes of gases 31 halogenoalkanes 464–5 macroscopic properties of 32–5
relationship between pressure and
fragmentation patterns 531–2 naming 432–3 temperature 34
free radical substitution 448–50 primary 440 relationship between pressure and volume
secondary 440 32–3
homolytic fission 449 tertiary 440 relationship between volume and
initiation 448–9 halogens/group 17 elements 98–100 temperature 33
propagation 449 atomic and ionic radii for 90 incomplete combustion 190
termination 449 displacement reactions of 100 independent variables 522
free radicals 174, 326, 449 formation during electrolysis 408 index of hydrogen deficiency (IHD) 524,
frequency factor 268 melting points of 99
frequency (waves) 62 and rate of nucleophilic substitution 469 525
fuel cells 415 reactions of 99 indicators 314, 349–53
fullerene 146–7
functional group isomers 443 with alkanes 447–51 pH range of 350–2
functional groups 424–6, 485 with alkenes 452 pKa of 352–3
variation in electron affinity 94 inductive reasoning 250
galvanic cells see voltaic cells heat 185 inert gases 85
gas chromatography−mass spectrometry heat energy 2 infinite dilution 223
helium 87 infrared spectroscopy 526–30
(GC−MS) 532 Hess’s law 196–202, 217 integration trace 534
gas constant 36, 268, 300 heterocyclic compounds 492 intermolecular forces 148–59
gas law equation 34 heterogeneous mixtures 6–7 internal energy 185
gases 1 heterolytic fission 466, 474 intramolecular forces 148
hexane 157, 158 inversely proportional relationships 520
calculations involving volumes of 27–39 high-resolution NMR spectra 538–40 iodine(I) chloride (ICl) 151
converting volumes of gases to number of homogeneous catalysts 265 ionic bonding 119–25
moles 28–31 homogeneous mixtures 6 and ionic crystals 122–3
formula for solving moles questions 31 homologous series 424–8 ionic compounds 119, 121, 392
ideal see ideal gases boiling points and 427 electrical conductivity of 124
real 27 homolytic fission 174 enthalpy changes when dissolved in water
geometrical isomers 152 Hund’s rule 71 222–4
germanium 68 hybridisation 176–9 formulas of 122
giant covalent structures 144–7 hydration enthalpies 224 melting points and boiling points of 123
giant ionic structures 123 hydrides, boiling points of 152–3 physical properties of 123–4, 159
Gibbs free energy 229 hydrocarbons 422 solubility 158
and equilibrium 232–3 boiling points 149–50
equilibrium constants and 300–1 naming 428–32 in non-polar solvents 124
gradient (slope) 242, 522, 523 see also alkanes; alkenes in water 124
graphene 146, 147 hydrochloric acid 320 volatility of 124
graphite 145–6 hydrogen 85, 126 ionic product constant (Kw) 316, 334
electrodes 408–9 emission spectra 63, 65 variation with temperature 336
graphs 519–24 ionisation energy of 72, 73 ionic radii 90–1
deriving quantities from 522–4 isotopes of 58 ionisation energy 72, 73–7, 92–3, 215
drawing 521–2 proton environments 534–5 of hydrogen 73
group 1 metals see alkali metals/group 1 reactions with alkenes 452 variation across a period 78–9
hydrogen bonding 152–4, 157 variation down a group 79, 86
metals hydrogen bromide ions 57, 119
group 17 elements see halogens/group 17 reaction with ethene 473–4 complex ions 108–9
reaction with propene 474–6 concentration of 42
elements hydrogen cyanide (HCN) 127, 340 formulas of 121–2
group 18 elements see noble gases hydrogen fluoride 129, 153 isoelectronic 119, 136
groups 85 hydrogen halides 452 shapes of 140–1, 164–6
iron 68, 79, 120, 393
variation in
electron affinity 94
electronegativity 95, 129
ionisation energy 86, 92
588
iron(III) chloride 356–7 mass 4 neutrons 56
isoelectronic ions 119, 136 conservation of 18–9 NF3 132
isomers 441–3 mass:charge ratio 61, 530 nitric acid 320, 326
of a molecule 11 nitrobenzene 479, 482–3
conformational 494–5 questions involving masses of substances nitrogen 79, 126
E/Z isomers 490–1, 492, 493 19–21 nitrogen oxides 102, 103, 325
functional group 443 see also relative atomic mass
geometrical 152 ozone depletion by 174, 175
see also cis-trans isomerism mass number 57 NO2– ion (nitrite) 136, 141, 171, 172
isotopes 58 mass spectrometer 60, 61 NO2+ ion (nitronium) 136
mass spectrometry 60–1, 530–3 NO3– ion (nitrate) 135, 171, 172
Ka see acid dissociation constant (Ka) Maxwell−Boltzmann distribution 249, 250 noble gases 85, 87, 129
Kb see base ionisation (dissociation) constant mechanisms of reactions 261–7 non-polar solvents, solubility of ionic solids
melting points 154, 160, 221, 222
(Kb) in 124
Kc see equilibrium constants (Kc) of alkali metals 97 non-spontaneous reactions 231
Kelvin temperature scale 29 of halogens 99 nuclear charge 90, 91, 95, 96
ketones 461 of ionic compounds 123
metal hydrides 370 trends across periods 78, 89, 92, 96
naming 436 metallic bonding 160–1 trends down groups 92
reduction of 481 metalloids 85 nuclear magnetic resonance (NMR)
kinetic energy of particles 2 metals 86
Kw see ionic product constant (Kw) alkali see alkali metals/group 1 metals spectroscopy 533–5, 536, 540–3
delocalised electrons in 160 nucleons 56
lanthanoide elements 85 reactions of acids with 312, 323 nucleophiles 463, 465, 466, 469
lattice enthalpies 216, 220–2, 224 transition see transition metals nucleophilic subsitution
lattice structure 123 methane (CH4) 126, 138, 140
Le Chatelier’s principle 282, 290, 349 bonding in 177 effect of the halogen of the rate of 469
combustion of 186–7 effect of the nucleophile on the rate of
effect of concentration 285 methanol 439 469
effect of pressure 283–4 methyl ethanoate 443 effect of the solvent on the rate of 471
effect of temperature 282–3, 290 methyl methanoate 443 mechanisms 466–72
lead bromide 390–1 methylbenzene 543 reactions 463, 464–5
Lewis acids and bases 310–1 2-methylbut-1-ene 431 see also SN1 mechanism; SN2 mechanism
Lewis structures 126, 132–7, 162–3 3-methylhexane 429 nucleus 56
ligands 108, 113–4 4-methylpent-2-yne 431
limiting reactants 24–6 2-methylpentane 429 octan-1-ol 158
line spectra 63, 64, 65 2-methylpropan-2-ol 439 octet rule 130–1
linear combination of atomic orbitals mixtures 6–7
molar mass 9 exceptions to 164
(LCAO) 167 molar volume 28 optical isomerism 495–8
linear shape 139, 142 molecular formula 14–7, 422
linked reactions 49 molecular ion 530 and plane-polarised light 497
liquid−vapour equilibrium 281 molecular mass 11 and ring compounds 496–7
liquids 1 molecular orbitals 167 orbitals 69–71, 167
molecular shapes 137–42, 164–6 in coloured complexes 110
see also solutions bent 139, 140, 141 order of reaction 252–3, 467
lithium 381 tetrahedral 138, 145 and rate expression from experimental
log scale 75 trigonal bipyramid 165 data 253–61
London (dispersion) forces 148, 149–50, trigonal planar 139 organic compounds
trigonal pyramidal 138, 141, 142 containing oxygen 434–9
155 molecularity 266, 467 determination of the number of multiple
in graphite 145 moles 9, 10, 19–21 bonds/rings in 524–6
between hexane and pentane molecules equation for solving moles questions oxidation 368, 372–3
157 involving solutions 44 oxidation numbers 112, 369–71
lone pairs of electrons 126, 140 monomers 455, 456, 457 naming compounds using 371–2
multiplicity 539 oxidation and reduction in terms of
magnesium 119, 507, 508 372–3
mass spectrum of 61 N2H4, shape of 141–2 of transition metals 106, 109, 370, 371
melting point of 160 neon 78 oxides
neutralisation 313 acidic 101, 102
magnesium chloride 45, 220, 221 neutrality 334 basic 101
magnesium fluoride 121, 122 of period 2 and period 3 elements 101–3
magnesium oxide 98, 101, 123, 218–9 oxidising agents (oxidants) 373, 402–3
magnetism 107 oxygen 119, 126, 173, 214
malleability 161 dissolved oxygen 382–3
margarine 453 electron configuration of 79
Markovnikov’s rule 475–6 reactions with alkali metals 98
INDEX 589
ozone (O3) 173, 179, 214 pKa 328–9 rate-determining step 262, 263, 264
delocalised system in 171 of an indicator 352–3 rate equation/rate expression 252–3
depletion by CFCs and NOx 174–5 from a titration curve 347
Lewis structures 170 from experimental data 253–60
resonance structures of 133, 170 pKb 330–1 experimental determination of 253
plane-polarised light, optical isomers and rate of reaction 241–5
paramagnetism 107 and catalysis 250
partial pressure 286, 300 497–8 and changes in colour 245
particles, number of 11–2 pOH 333–4 and concentration 247, 252
parts per million (ppm) 41, 42 point of inflexion 349 and decrease in mass 243
Pauli exclusion principle 70 polar molecules 143, 150–2 definition of 244–5
pentan-3-one 534 polarimeters 498 equilibrium and 279
pentane 157 polarity 128–30 equilibrium constants and 291
percentage composition of compounds polyethene 455, 457 factors affecting 247
polymerisation, addition 455–7 and pressure 247
12–4 polymers 457 and surface area of solid reactants 248
percentage errors 511, 512 polypropene 456 and temperature 248–9
percentage uncertainties 516 polythene (polyethene) 455, 457 reaction pathways 484–7
percentage yield, of chemical reactions polyunsaturated oil 453 reaction quotient (Q) 287–8, 298
positive inductive effect 470, 475 reaction time 509
23–4 potassium 381 real gases 27
periodic table 85–8 redox equations 374–80
ionisation energies of 74–5 redox reactions 100
and atomic radii 88–9 potassium chloride 392 and the activity series 381
and electron configurations 87–8 potential energy diagrams (profiles) predicting 381, 401
and first ionisation energy 92–3 and standard electrode potentials 401
groups see groups for a reversible reaction 290, 291 see also oxidation; reduction
and ionic radii 90–1 showing activation energy for an redox titrations 382–3
periods see periods exothermic reaction 246 reducing agents (reductants) 373, 402–3
and physical properties 88–103 and stability 213 reduction 368, 481–3
periods 85 for a two-step reaction 262 of aldehydes 481
variation in precision 510, 512 of carboxylic acids 482
pressure of ketones 481
electron affinity 94–5 and equilibrium 283–4, 292 in terms of oxidation numbers 372–3
electronegativity 96, 129 partial pressure 286, 300 see also redox reactions
ionic radii 90–1 primary alcohols 458–61 refrigeration 2
ionisation energy 78–9, 92–3 primary halogenalkanes 469–70 relative atomic mass 8, 59–60, 149
permanent dipole−permanent dipole SN2 mechanism 466–7 relative formula mass 8
interactions 150 principal quantum number 62 relative molecular mass 8–9, 149, 155
PF5 (phosphorus(V) fluoride) 164–5 propan-1-ol 528, 543 repeating units 456
pH 324 propan-2-ol 439 resonance 170–3
of bases 333–4 propane 150–1 resonance hybrids 170
calculation of 317–8 propanoic acid 443, 531 resonance structures 133–4, 170
[H+(aq)] from 316 propene, reaction with HBr 474–6 retrosynthesis 485
indicator range of 350–2 proportional relationships 520, 521 reversible reactions 278
measurement of 314–5 protons 56, 533 see also equilibrium
neutral 334 PVC (polyvinylchloride) 456 ring compounds, optical isomerism and
of strong acids/bases at different
temperatures 337 qualitative data 7, 507 496–7
of a weak acid 331–3 quantitative data 7, 507 rounding values in calculations 516
pH curves 341–53 quantitative electrolysis 413–5
pH meter 315 salt bridge 387
pH scale 314–9 current 413 salt hydrolysis 354–7
phase equilibrium 281 ion charge 413 salts 312
phenylamine (aniline) 330, 482, 483 time of electrolysis 413
phosphate(V) ion 163 quartz 147 from acids and bases 314, 354–6
phosphorus 95 saturated compounds 431, 451
oxides of 102 racemic mixtures 499 scandium 104
phosphorus(V) fluoride (PF5) 164–5 radio frequency 533 scientific law 38
photons 63, 72–3 radioactivity 58 sea water 7
physical properties 3, 88–103 radioisotopes 279 second electron affinity 215
see also boiling points; melting points random uncertainties 508, 509 second ionisation energy 74, 215
pi (π) bonding 114, 167–9 rate constant (k) 252 second law of thermodynamics 228
pi (π) delocalised systems 146
effect of catalysts on 271
590 units 260
second-order reactions 259–60 standard enthalpy change 187, 188 rate of reaction and 248–9
secondary alcohols 461 of atomisation 215 and spontaneity of a reaction 231–2
selenium 68 of combustion 188 variation of Kw with 336
SF4 (sulfur(IV) fluoride) 165 of formation 203–6 temperature scale, absolute/Kelvin 29
SF6 (sulfur(VI) fluoride) 164 of a reaction 188 tertiary alcohols 462
shapes of ions see ions, shapes of of vaporisation 207 tertiary halogenoalkenes 468–9, 470
shapes of molecules see molecular shapes tetrahedral shape 138, 145
shells, electron 62 standard entropy 225 tetramethylsilane (TMS) 536
standard free energy change 229–30, 405 thalidomide 500
see also energy levels; subshells theoretical yield, of chemical reactions 23
sigma bonds 167–9 electrolysis and 412 theory 38
significant figures 513, 514, 515 standard hydrogen electrode 394 tin 86
silicon 76, 77, 95 standard solution 43 titration curves 341–53
silicon dioxide 147 standard temperature and pressure (STP) 28 pKa from 347
single bonds 127, 128, 170 state symbols 5, 203 titrations 43, 324
single crystal X-ray crystallography 544 states of matter 1 back titrations 47–8
skeletal formulas 423 stereoisomerism 489–501 end point of 349
SN1 mechanism 265–7, 468–9, 470 transition elements see transition metals
SN2 mechanism 265–7, 466–7, 469–70 cis-trans isomerism 489–94 transition metals 104–5
conformation isomers 494–5 catalytic ability 109
rate-determining step 472 diastereomers 499 complex ions 108–9, 110–4
SO2, Lewis structures 134 optical isomerism 495–8
SO42– (sulfate ion) 370 stereospecific reaction 467 d orbitals 112
sodium 89, 160 steric effects 469–70 ligands 113–4
sodium chloride 122–3, 158, 221, 356 STP (standard temperature and pressure) 28 oxidation number 109, 112
strong acids 320 electron configurations of 79, 105
Born−Haber cycle 216–8 calculation of pH of 317 ionisation of 105
sodium hydroxide 321 conjugate base 322 ions of 120, 122
sodium oxide 101 distinguishing weak and 323 magnetic properties of 107
solids 1 titration against a strong base 341–3 oxidation numbers of 106, 109, 370, 371
solubility 156–8 titration against a weak base 347–8 properties of 104–5
solute 39 strong bases 321 transition state 466, 467
solutions calculation of pH of 318 1,1,2-trichloroethane 538–9
conjugate acid 322 trigonal bipyramid shape 165
calculations involving 39–51 titration against a strong acid 341–3 trigonal planar shape 139
concentration of 40, 41–2 structural formula 422–3 trigonal pyramidal shape 138, 141, 142
definition of 39 structural isomerism 441, 493 3,3,4-trimethylhexane 430
enthalpy changes in 191–5 sub-energy levels 67 2,6,6-trimethyloctane 430
equation for solving moles questions sublimation 1 triple bonds 126, 127, 168, 170
involving 44 subshells 67, 70, 77, 87
solvents 39 transition metals 104, 106, 111 ultraviolet light 173–5, 214
aprotic 471 substitution reactions 448, 454, 464 uncertainties
effect on the rate of nucleophilic sulfate(VI) ion 163
substitution 471 sulfur 87 in mean values 513
protic 471 oxides of 102 in measurements 507–14
specific heat capacity 188–9 sulfur dioxide 139, 325 propagating in calculations 516–9
spectrochemical series 113 formal charges 162–3 quoting values with 513
spectroscopes 63 methods of reducing emissions of 327 see also errors
spectroscopy, using combined techniqes to oxidation number 370 Universal indicator 314, 315, 323, 350
sulfuric acid 320 Universe, entroy of the 228–9
determine molecular structures 545–7 sulfur(IV) fluoride 165 unsaturated compounds 431, 451
spin sulfur(VI) fluoride 164 unsaturated fats 453
surroundings 185
electron 70 system 185 valence shell electron pair repulsion theory
hydrogen nucleus 533 systematic errors 510, 511–2 (VSEPR) 137–42
spin−spin splitting/coupling 539, 540
spontaneous reactions 228 tangents 242, 523 van der Waals’ forces 148
stability 187 temperature 185 vaporisation, standard enthalpy change of
standard cell potentials 393, 394, 405
standard deviation 510 and the energy of particles in a gas 248 207
standard electrode potentials 393–406 and equilibrium 282–3, 292 variables 522
measurement of 395–400 and equilibrium constant 289–90 vibrational energy 526
to predict pH and 337 volatility 124
voltage 389
feasibility of a redox reaction 401
product of electrolysis 407–8
INDEX 591
voltaic cells 386–90, 392
anodes 387
cathodes 388
cell notation for 388–9
voltmeters 395
VSEPR (valence shell electron pair
repulsion) theory 137–42
water 140, 311
boiling point of 153
bonding in 126
dissociation of 316
electrolysis of 412
reactions with alkali metals 98
reactions with alkenes 452–3
solubility in 157
of ionic substances 124
water of crystallisation 45–7
wavelength 73
wavenumbers 526
weak acids 320–1
titration against a strong base 345–7
titration against a weak base 348–9
weak bases 322
calculating pH or pOH for 334–5
titration against a weak acid 348–9
white light 63, 111
Winkler titration method 383
X-ray crystallography 139, 544
xenon(IV) fluoride (XeF4) 164, 166
yield of a chemical reaction 23–4
zero-order reactions 258, 260
zinc 104, 110, 386
592
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Chemistry for the IB Diploma Coursebook 2nd ed - Steve Owen (CUP, 2014)
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