IB Chemistry Study Guide: 2014 Edition: Oxford IB Diploma Program

210 Pages·2014·13.524 MB·English

by  Geoff Neuss

O X F O R D I B S T U D Y G U I D E S
Geofrey Neuss
Ch e m i s try
FO R TH E I B D I PLO M A
2014 edition
2
3
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Acknowledgements
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Introduction and acknowledgements
This book replaces the very successul rst and second I have stuck rigorously to the syllabus to produce this
editions. Like earlier editions it is written specically study guide which contains all the necessary subject
or students studying Chemistry or the International content required or the examination in one easily
Baccalaureate Diploma although many students ollowing accessible and compact ormat. IB Chemistry is, o
their own national systems will also nd it helpul. It course, about much more than the nal examination
comprehensively covers the new programme that will and you are also encouraged to read widely around the
be examined rom 2016 onwards. All the inormation subject to urther your knowledge and understanding
required or each topic is set out in separate boxes with as well as your enjoyment o chemistry. Many more
clear titles that ollow aithully the layout o the syllabus. multiple choice tests and short answer questions, as
The rst 11 topics cover both the Core content needed well as a blog and much more background inormation
by all students and the Additional Higher Level material (e.g. videos) about each topic and option can be ound
under the main topic headings. The dierence between on my website or IB Diploma Chemistry teachers at
the two levels is clearly distinguished. Both Higher Level www.thinkib.net/chemistry.
and Standard Level students must study one o the our I have been ortunate at Atlantic College to teach many
options and each option stands in its own right even i highly motivated and gited students who have oten
this has meant repeating small amounts o some o the challenged me with searching questions. During my
material. Worked examples are included where they are association with the International Baccalaureate, the
appropriate and at the end o each main topic there are European Baccalaureate and the United World Colleges
both multiple choice and short-answer practice questions. I have been privileged to meet, work alongside and
Similarly at the end o each option I have included short- exchange ideas with many excellent Chemistry teachers
answer questions. Many o these questions are taken who exude a real enthusiasm or their subject. Many
rom past IB examination papers and I would like to thank o these have infuenced me greatly  in particular,
the International Baccalaureate Organization or giving John Devonshire, a ellow teacher at Atlantic College
me permission to use them. The remaining questions and Jacques Furnemont, an Inspector o Chemistry
are written to the same IB standard specically or this in Belgium. I value greatly their advice, opinions and
book. Worked answers to the questions are provided. knowledge. I would also like to pay tribute to two ormer
These answers are not necessarily ull 'model' answers Chie Examiners or the IB, Ron Ragsdale and Arden Zipp.
but they do contain all the inormation needed to score The high regard with which IB Chemistry is held today
each possible mark. I have included one chapter on the owes much to both o them.
underlying philosophy o the course to explain and give
Finally I should like to thank my wie Chris and my riend
examples o the Essential Ideas, the Nature o Science,
and colleague John or their patience and unstinting
the International Dimension, and Utilization. To help you,
support throughout.
the student, gain the highest grade possible the nal
chapter is devoted to giving you advice on how to study
and prepare or the nal examination. It also advises Dr Georey Neuss
you on how to excel at the internally assessed practical
component o the course. For those who opt or Chemistry
as the subject or their Extended Essay it gives advice
and guidance on how to choose the topic and write your
Essay. A comprehensive Extended Essay checklist is
included to help you gain bonus points towards your IB
Diploma.
IntroductIon and acknowledgements iii
Contents
(Italics denote topics which are exclusively Higher Level.)
1 stoIcHIometrIc relatIonsHIPs Hybridization (2) 35
(IB toPIc 1) Multiple choice questions 36
Short answer questions 37
Particulate nature o matter 1
5 energetIcs / tHermocHemIstrY
The mole concept and chemical ormulas 2
(IB toPIc 5 and toPIc 15)
Chemical reactions and equations 3
Mass and gaseous volume relationships 4 Measuring enthalpy changes 38
Molar volume o a gas and calculations 5 H calculations 39
Titration and atom economy 6 Hess Law and standard enthalpy changes 40
Multiple choice questions 7 Bond enthalpies 41
Short answer questions 8 Energy cycles 42
2 atomIc structure Entropy and spontaneity 43
(IB toPIc 2 and toPIc 12) Spontaneity o a reaction / Multiple choice questions 44
Multiple choice questions (continued) 45
The nuclear atom 9
Short answer questions 46
Mass spectrometer and relative atomic mass 10
6 cHemIcal kInetIcs
Emission spectra 11
(IB toPIc 6 and toPIc 16)
Electronic conguration 12
Evidence rom ionization energies 13 Rates o reaction and collision theory 47
Multiple choice questions 14 Factors afecting the rate o reaction 48
Short answer questions 15 Rate expression and order o reaction 49
3 PerIodIcItY (IB toPIc 3 and Reaction mechanisms and activation energy 50
toPIc 13) Multiple choice questions 51
Multiple choice questions (continued) 52
The periodic table 16
Short answer questions 53
Periodic trends (1) 17
7 eQuIlIBrIum
Periodic trends (2) 18
(IB toPIc 7 and toPIc 17)
Periodic trends (3) 19
The transition metals 20 The equilibrium law 54
Transition metal complex ions 21 Le Chateliers principle and actors afecting the
Colour o transition metal complexes 22 position o equilibrium 55
Multiple choice questions 23 Equilibrium calculations 56
Short answer questions 24 Multiple choice questions 57
Short answer questions 58
4 cHemIcal BondIng and
8 acIds and Bases
structure (IB toPIc 4 and toPIc 14)
(IB toPIc 8 and toPIc 18)
Ionic bonding 25
Covalent bonding 26 Theories and properties o acids and bases 59
Shapes o simple molecules and ions 27 The pH scale 60
Resonance hybrids and allotropes o carbon 28 Strong and weak acids and bases and simple pH
Intermolecular orces 29 calculations 61
Physical properties related to bonding type 30 Acid deposition 62
Metals and alloys 31 Lewis acids and bases 63
Molecular orbitals 32 Calculations involving pH, pOH and pK 64
w
Oxygen and ozone 33 Calculations with weak acids and bases 65
Hybridization (1) 34 Salt hydrolysis and bufer solutions 66
iv contents
Titration curves and indicators 67 Spectroscopic identication o organic
Multiple choice questions 68 compounds  IR and 1H NMR 103
Short answer questions 69 Nuclear magnetic resonance (NMR)
spectroscopy 104
9 redoX Processes
Applications o 1H NMR spectroscopy 105
(IB toPIc 9 and toPIc 19)
Combination o diferent analytical techniques to
Redox reactions (1) 70 determine structure 106
Redox reactions (2) 71 X-ray crystallography / Multiple choice questions 107
Activity series and Winkler method 72 Multiple choice questions continued 108
Electrochemical 73 Short answer questions 109
Electrolysis 74 12 oPtIon a  materIals
Electroplating and standard electrode potentials 75
Introduction to materials science 110
Spontaneity o electrochemical reactions 76
Principles o extraction o metals rom their ores 111
Multiple choice questions 77
Faraday calculations and properties and analysis
Short answer questions 78 o alloys 112
10 organIc cHemIstrY Catalysts 113
(IB toPIc 1o and toPIc 20) Liquid crystals (1) 114
Liquid crystals (2) 115
Fundamentals o organic chemistry 79
Polymers 116
Common classes o organic compounds 80
Nanotechnology 117
Structural ormulas 81
Environmental impact  plastics 118
Structural isomers 82
Superconducting metals and X-ray crystallography 119
3-D models o structural ormulas 83
Condensation polymers 120
Properties o diferent homologous series 84
Environmental impact  heavy metals (1) 121
Alkanes 85
Environmental impact  heavy metals (2) 122
Alkenes 86
Short answer questions 123
Alcohols 87
13 oPtIon B  BIocHemIstrY
Substitution and condensation reactions 88
Nucleophilic substitution 89 Introduction to biochemistry 124
Electrophilic addition reactions (1) 90 Structure o proteins 125
Electrophilic addition reactions (2) 91 Analysis o proteins 126
Electrophilic substitution reactions and reduction Enzymes 127
reactions 92 Lipids (1) 128
Synthetic routes 93 Lipids (2) 129
Stereoisomerism (1) 94 Carbohydrates 130
Stereoisomerism (2) 95 Vitamins 131
Stereoisomerism (3) 96 Biochemistry and the environment 132
Multiple choice questions 97 Proteins and enzymes (1) 133
Short answer questions 98 Proteins and enzymes (2) 134
11 measurement, data Nucleic acids 135
ProcessIng and analYsIs The genetic code 136
(IB toPIc 11 and toPIc 21) Biological pigments (1) 137
Biological pigments (2) 138
Uncertainty and error in measurement 99
Stereochemistry in biomolecules 139
Uncertainty in calculated results and graphical
techniques 100 Short answer questions 140
Analytical techniques 101 14 oPtIon c  energY
Spectroscopic identication o organic Energy sources 141
compounds  MS 102
Fossil uels (1) 142
contents v
16 UnDeRLYInG PHILosoPHY
Fossil uels (2) 143
Nuclear usion and nuclear ssion (1) 144 Introduction 167
Nuclear usion and nuclear ssion (2) 145 Essential ideas (1) 168
Solar energy 146 Essential ideas (2) 169
Environmental impact  global warming 147 Nature o Science (1) 170
Electrochemistry, rechargeable batteries and Nature o Science (2) 171
uel cells (1) 148
International-mindedness 172
Electrochemistry, rechargeable batteries and
Utilization 173
uel cells (2) 149
17 oBtAInInG A HIGH FInAL
Nuclear usion and nuclear fssion (3) 150
GRADe
Nuclear usion and nuclear fssion (4) 151
Photovoltaic cells and dye-sensitized Study methods 174
solar cells (DSSC) 152 The nal examinations 175
Short answer questions 153 Command terms 176
15 oPtIon D  MeDICInAL Internal Assessment (1) 177
CHeMIstRY Internal Assessment (2) 178
Extended Essays (1) 179
Pharmaceutical products and drug action 154
Extended Essays (2) 180
Aspirin 155
Extended Essays (3) 181
Penicillin 156
Opiates 157
ANSWERS TO QUESTIONS 182
pH regulation o the stomach 158
ORIGIN OF INDIVIDUAL QUESTIONS 190
Bufer solutions 159
INDEX 191
Antiviral medications 160
PERIODIC TABLE FOR USE WITH THE IB 196
Environmental impact o some medications 161
Taxol  a chiral auxiliary case study 162
Nuclear medicine 163
Drug detection and analysis (1) 164
Drug detection and analysis (2) 165
Short answer questions 166
vi Contents
1 S TO I C H I O M E T R I C R E L AT I O N S H I PS
Particulate nature of matter
ELEMENTS COMPOUNDS
All substances are made up o one or more elements. An Some substances are made up o a single element, although
element cannot be broken down by any chemical process into there may be more than one atom o the element in a
simpler substances. There are just over 100 known elements. particle o the substance. For example, oxygen is diatomic,
The smallest part o an element is called an atom. that is, a molecule o oxygen contains two oxygen atoms and
Names o the frst 20 elements has the ormula O2. A compound contains more than one
element combined chemically in a fxed ratio. For example,
Atomic Name Symbol Relative
a molecule o water contains two hydrogen atoms and one
Number atomic mass
oxygen atom. It has the ormula HO. Water is a compound
1 hydrogen H 1.01 not an element because it can be b2roken down chemically
2 helium He 4.00 into its constituent elements: hydrogen and oxygen.
3 lithium Li 6.94 Compounds have dierent chemical and physical properties
4 beryllium Be 9.01 rom their component elements.
5 boron B 10.81
6 carbon C 12.01
7 nitrogen N 14.01 MIXTURES
8 oxygen O 16.00
The components o a mixture may be elements or
9 uorine F 19.00
compounds. These components are not chemically bonded
10 neon Ne 20.18
together. Because they are not chemically combined, the
11 sodium Na 22.99
components o a mixture retain their individual properties.
12 magnesium Mg 24.31
All the components o a mixture may be in the same phase,
13 aluminium Al 26.98
in which case the mixture is said to be homogeneous. Air
14 silicon Si 28.09
is an example o a gaseous homogenous mixture. I the
15 phosphorus P 30.97
components o a mixture are in dierent phases the mixture
16 sulur S 32.07
is said to be heterogeneous. There is a physical boundary
17 chlorine Cl 35.45
between two phases. A solid and a liquid is an example o a
18 argon Ar 39.95
two-phase system. It is possible to have a single state but two
19 potassium K 39.10 phases. For example, two immiscible liquids such as oil and
20 calcium Ca 40.08 water orm a heterogeneous mixture.
STATES OF MATTER
sublimation
SOLID STATE LIqUID STATE GASEOUS STATE
 Fixed shape  Fixed volume  Widely spaced particles that
 Fixed volume  Takes up shape o container completely fll container
 Particles held together by  Particles held closely together  Pressure o the gas due to
gaseous particles colliding with
intermolecular orces in a by intermolecular orces
the walls o the container
fxed position  Particles have translational
 Particles can vibrate about a velocity so diusion can occur boiling/  Intermolecular orces between
particles negligible
fxed point but do not have  As heat is supplied the liquid vaporization/
translational velocity melting particles move aster. Some evaporation  Volume occupied by
 As heat is supplied at a certain particles move aster than others molecules themselves
reezing condensation negligible compared with total
temperature the vibration and escape rom the surace
volume o gas
is sufcient to overcome o the liquid to orm a vapour.
the attractive orces holding Once the pressure o the vapour  Particles moving with rapid,
the solid together and the is equal to the pressure above random motion so diusion
solidmelts the liquid the liquid boils. can occur
deposition
STOICHIOMETRIC RELATIONSHIPS 1
The mole concept and chemical formulas
MOLE CONCEPT AND AVOGADROS CONSTANT
A single atom of an element has an extremely small mass. For example, an atom of carbon-12 has a mass of 1.993  10-23 g.
This is far too small to weigh. A more convenient amount to weigh is 12.00 g. 12.00 g of carbon-12 contains 6.02  1023 atoms of
carbon-12. This number is known as Avogadros constant (N or L).
A
Chemists measure amounts of substances in moles. A mole is the amount of substance that contains L particles of that substance.
The mass of one mole of any substance is known as the molar mass and has the symbol M. For example, hydrogen atoms have _ 1_
12
of the mass of carbon-12 atoms so a mole of hydrogen atoms contains 6.02  1023 hydrogen atoms and has a mass of 1.01 g. In
reality elements are made up of a mixture of isotopes.
The relative atomic mass of an element A is the weighted mean of all the naturally occurring isotopes of the element relative to
r
carbon-12. This explains why the relative atomic masses given for the elements on page 1 are not whole numbers. The units of molar
mass are g mol-1 but relative molar masses M have no units. For molecules relative molecular mass is used. For example, the M of
r r
glucose, CH O = (6  12.01) + (12  1.01) + (6  16.00) = 180.18. For ionic compounds the term relative formula mass is used.
6 12 6
Be careful to distinguish between the words mole and molecule. A molecule of hydrogen gas contains two atoms of hydrogen and has
the formula H. A mole of hydrogen gas contains 6.02  1023 hydrogen molecules made up of two moles (1.20  1024) of hydrogen atoms.
2
FORMULAS OF COMPOUNDS
Compounds can be described by different chemical formulas. substance. It can be obtained from the empirical formula if the
Empirical formula (literally the formula obtained by experiment) molar mass of the compound is also known.
This shows the simplest whole number ratio of atoms of each Methanal CH2O (Mr = 30), ethanoic acid C2H4O2 (Mr = 60)
element in a particle of the substance. It can be obtained by and glucose C6H12O6 (Mr = 180) are different substances with
either knowing the mass of each element in the compound or different molecular formulas but all with the same empirical
from the percentage composition by mass of the compound. The formula CH2O. Note that subscripts are used to show the
percentage composition can be converted directly into mass by number of atoms of each element in the compound.
assuming 100 g of the compound are taken. Structural formula
Example: A compound contains 40.00% carbon, 6.73% This shows the arrangement of atoms and bonds within a
hydrogen and 53.27% oxygen by mass, determine the molecule and is particularly useful in organic chemistry.
empirical formula. The three different formulas can be illustrated using ethene:
H H
Amount / mol Ratio
C C
C 40.00/12.01 = 3.33 1
H 6.73/1.01 = 6.66 2 Empirical formula CH C H H H
= CHO 2 2 4
O 53.27/16.00 = 3.33 1 2 (can also be written
Molecular formula HC=CH)
2 2
For molecules this is much more useful as it shows the empirical molecular structural
actual number of atoms of each element in a molecule of the formula formula formula
EXPERIMENTAL DETERMINATION OF AN EMPIRICAL FORMULA
The empirical formula of magnesium oxide can be determined simply in the laboratory. A coil of magnesium ribbon about
10cm long is placed in a pre-weighed crucible and its mass recorded. The crucible is placed on a clay triangle and heated
strongly. When the magnesium ribbon starts to burn the lid is lifted slightly to allow more air to enter and the heating is
continued until all the magnesium has burned. After cooling the crucible, its lid and its contents are reweighed.
Table of typical raw quantitative data:
Mass / g ( 0.001 g) lid
Mass of crucible + lid 30.911
Mass of crucible + lid + magnesium 31.037
crucible
Mass of crucible + lid + magnesium oxide 31.106
coil of magnesium
Mass of magnesium = 31.037 - 30.911 = 0.126 g
Mass of magnesium oxide = 31.106 - 30.911 = 0.195 g
Mass of oxygen combining with magnesium = 0.195 - 0.126 = 0.069 g
Amount of magnesium = _0.126 = 5.2  103 mol
24.31
Bunsen burner
Amount of oxygen = _ 0.069 = 4.3  103 mol
16.00
Ratio of Mg to O = _ 5.2 _ 103
4.3  103 = 1.2 to 1
Convert to whole number ratio = 6 : 5
Empirical formula of magnesium oxide as determined by this experiment is MgO .
6 5
2 STOICHIOMETRIC RELATIONSHIPS
Chemical reactions and euations
PROPERTIES OF CHEMICAL REACTIONS
In all chemical reactions:
 new substances are ormed
 bonds in the reactants are broken and bonds in the products are ormed resulting in an energy change between the reacting
system and its surroundings
 there is a fxed relationship between the number o particles o reactants and products resulting in no overall change in mass 
this is known as the stoichiometry o the reaction.
CHEMICAL EqUATIONS
Chemical reactions can be represented by chemical equations. Reactants are written on the let-hand side and products on the
right-hand side. The number o moles o each element must be the same on both sides in a balanced chemical equation, e.g. the
reaction o nitric acid (one o the acids present in acid rain) with calcium carbonate (the main constituent o marble statues).
CaCO (s) + 2HNO (aq)  Ca(NO ) (aq) + CO (g) + H O(l)
3 3 3 2 2 2
calcium carbonate nitric acid calcium nitrate carbon dioxide water

REACTANTS PRODUCTS
STATE SyMbOLS
Because the physical state that the reactants and products are in can aect both the rate o the
reaction and the overall energy change it is good practice to include the state symbols in the
equation.
(s) - solid (l) - liquid (g) - gas (aq) - in aqueous solution
 OR 
A single arrow  is used i the reaction goes to completion. Sometimes the reaction conditions are
written on the arrow:
Ni catalyst, 180 C
e.g. C H(g) + H(g)  C H(g)
2 4 2 2 6
Reversible arrows are used or reactions where both the reactants and products are present in the
equilibrium mixture:
Fe(s), 550 C
e.g. 3H(g) + N(g) 2NH(g)
2 2 3
250 Atm
COEFFICIENTS AND MOLAR RATIO
The coefcient reers to the number in ront o each reactant and product in the equation. The
coefcients give inormation on the molar ratio. In the frst example above, two moles o nitric acid
react with one mole o calcium carbonate to produce one mole o calcium nitrate, one mole o carbon
dioxide and one mole o water. In the reaction between hydrogen and nitrogen above, three moles o
hydrogen gas react with one mole o nitrogen gas to produce two moles o ammonia gas.
IONIC EqUATIONS
Because ionic compounds are completely dissociated in solution it is sometimes better to use ionic equations. For example, when
silver nitrate solution is added to sodium chloride solution a precipitate o silver chloride is ormed.
Ag+(aq) + NO -(aq) + Na+(aq) + Cl-(aq)  AgCl(s) + Na+(aq) + NO -(aq)
3 3
Na+(aq) and NO -(aq) are spectator ions and do not take part in the reaction. So the ionic equation becomes:
3
Ag+(aq) + Cl-(aq)  AgCl(s)
From this we can deduce that any soluble silver salt will react with any soluble chloride to orm a precipitate o silver chloride.
STOICHIOMETRIC RELATIONSHIPS 3
Mass and gaseous volume relationships
SOLIDS SOLUTIONS
Normally measured by weighing to obtain Volume is usually used for solutions.
the mass. 1.000 litre = 1.000 dm3 = 1000 cm3
1.000 kg = 1000 g Concentration is the amount of solute
When weighing a substance the mass (dissolved substance) in a known volume of
should be recorded to show the accuracy of solution (solute plus solvent). It is expressed
the balance. For example, exactly 16 g of a MEASUREMENT OF either in g dm-3, or, more usually in mol dm-3.
substance would be recorded as MOLAR qUANTITIES For very dilute solutions it is also sometimes
16.00 g on a balance weighing to + or expressed in parts per million, ppm. A
In the laboratory moles can
- 0.01 g but as 16.000 g on a balance solution of known concentration is known as
conveniently be measured
weighing to + or - 0.001 g. a standard solution.
using either mass or volume
depending on the substances To prepare a 1.00 mol dm-3 solution of
sodium hydroxide dissolve 40.00 g of solid
involved.
sodium hydroxide in distilled water and
then make the total volume up to 1.00 dm3.
Concentration is often represented by
LIqUIDS
square brackets, e.g.
Pure liquids may be weighed or the volume
recorded. GASES [NaOH(aq)] = 1.00 mol dm-3
The density of the liquid = _ v_om_lua_ms_s e_ and is Mass or volume may be used A 25.0 cm3 sample of this solution contains
usually expressed in g cm-3. for gases. 1.00  _ 12_05_0.00_ = 2.50  102 mol of NaOH
CHANGING THE VARIAbLES FOR A FIXED MASS OF GAS
P  _ V1_ (or PV = constant) P  T (or _ PT_ = constant) V  T (or _ VT_ = constant)
At constant temperature: At constant volume: increasing the At constant pressure: at higher
as the volume decreases the temperature increases the average temperatures the particles have a greater
concentration of the particles kinetic energy so the force with which average velocity so individual particles will
increases, resulting in more the particles collide with the container collide with the container walls with greater
collisions with the container walls increases. Hence pressure force. To keep the pressure constant there
walls. This increase in pressure increases and is directly proportional must be fewer collisions per unit area so
is inversely proportional to to the absolute temperature, i.e. the volume of the gas must increase. The
the volume, i.e. doubling the doubling the absolute temperature increase in volume is directly proportional to
pressure halves the volume. doubles the pressure. the absolute temperature, i.e. doubling the
absolute temperature doubles the volume.
IDEAL GAS EqUATION
e
r The different variables for a gas are
u
ss all related by the ideal gas equation. me
pre PV = nRT olu
v
P = pressure in Pa (N m-2)
volume (1 atm = 1.013  105 Pa)
T = absolute temperature in K 0 273 373 / K
V = volume in m3 273 0 100 / C
ure (1 cm3 = 1  106 m3) temperature
s
es n = number of moles Extrapolating the graph to zero volume gives
r
p
R = gas constant = 8.314 J K1 mol1 the value for absolute zero.
1 / volume
UNITS REAL GASES
The gas constant can be expressed in different units but it is An ideal gas exactly obeys the gas laws. Real gases do have
easier to use SI units. some attractive forces between the particles and the particles
R = _ PnVT = _N mmo-l2 _  K m 3 = N m mol1 K-1 tthhee mlaswelsv. eIsf tdhoe oyc dciudp tyh seoym coe uslpda cnee vsoer t hcoeny ddeon nseo tin etxoa lcitqluy iodbs.e y
= J K-1 mol-1 A gas behaves most like an ideal gas at high temperatures and
low pressures.
4 STOICHIOMETRIC RELATIONSHIPS