AQA GCSE (9–1)
Chemistry AQA GCSE (9–1) Chemistry Student Book helps to develop the skills, knowledge and confidence needed for the new GCSE course. This title has been approved by AQA.
supports effective learning, with straightforward explanations and colour coding on every page to show the level of challenge develops the maths skills you need for each topic, with Maths skills pages and activities consolidates understanding of all required practicals, helping you prepare for the indirect assessment helps you practise applying, interpreting and evaluating scientific ideas and data gives step-by-step guidance on how to improve answering questions, with annotated student answers
AQA GCSE (9–1) Chemistry
AQA GCSE (9–1) Chemistry Student Book:
covers both foundation and higher tier courses, with higher tier only content clearly labelled.
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AQA GCSE (9–1) Biology Student Book 978-0-00-815875-0
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AQA GCSE (9–1)
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AQA GCSE (9–1)
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*only the student book is approved by AQA and other resources have not been entered into the approval process.
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You can use this book if you are studying Combined Science: Trilogy
Contents
you will need to master all of the ideas and concepts on these pages you will need to master some of the ideas and concepts on these pages.
How to use this book
6
Chapter 1 Atomic structure and the periodic table 12
1.1
Elements and compounds Atoms, formulae and equations 1.3 Mixtures 1.4 Changing ideas about atoms 1.5 Modelling the atom 1.6 Relating charges and masses 1.7 Sub-atomic particles 1.8 Electronic structure 1.9 The periodic table 1.10 Developing the periodic table 1.11 Comparing metals and non-metals 1.12 Metals and non-metals 1.13 Key concept: The outer electrons 1.14 Exploring Group 0 1.15 Exploring Group 1 1.16 Exploring Group 7 1.17 Reaction trends and predicting reactions 1.18 Transition metals 1.19 Maths skills: Standard form and making estimates 1.2
14 16 18 20 22 24 26 28 30 32 34 36 38 40 42 44 46 48 50
Chapter 2 Structure, bonding and the properties of matter 56
2.1
Chemical bonds Ionic bonding 2.3 Ionic compounds 2.4 Covalent bonding 2.5 Metallic bonding 2.6 Three states of matter 2.7 Properties of ionic compounds 2.8 Properties of small molecules 2.9 Polymer structures 2.10 Giant covalent structures 2.11 Properties of metals and alloys 2.12 Diamond 2.13 Graphite 2.14 Graphene and fullerenes 2.15 Nanoparticles, their properties and uses 2.16 Key concept: Sizes of particles and orders of magnitude 2.17 Maths skills: Visualise and represent 2D and 3D shapes 2.2
58 60 62 64 66 68 70 72 74 76 78 80 82 84 86
90
Key concept: Conservation of mass and balanced equations 3.2 Relative formula mass 3.1
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98 100
102 104 106 108 110 112 114 116 118 120 122 124
Chapter 4 Chemical changes 130
88
Chapter 3 Chemical quantities and calculations 96
3.3 Mass changes when gases are in reactions 3.4 Chemical measurements and uncertainty 3.5 Moles 3.6 Amounts of substances in equations 3.7 Using moles to balance equations 3.8 Concentration of solutions 3.9 Key concept: Percentage yield 3.10 Atom economy 3.11 Using concentrations of solutions 3.12 Amounts of substance in volumes of gases 3.13 Key concept: Amounts in Chemistry 3.14 Maths skills: Change the subject of an equation 4.1
Metal oxides Reactivity series 4.3 Extraction of metals 4.4 Oxidation and reduction in terms of electrons 4.5 Reaction of metals with acids 4.6 Neutralisation of acids and salt production 4.7 Soluble salts 4.8 Required practical: Preparing a pure, dry sample of a soluble salt from an insoluble oxide or carbonate 4.9 pH and neutralisation 4.10 Required practical: Finding the reacting volumes of solutions of acid and alkali by titration 4.11 Strong and weak acids 4.12 The process of electrolysis 4.13 Electrolysis of molten ionic compounds 4.14 Using electrolysis to extract metals 4.15 Electrolysis of aqueous solutions 4.16 Required practical: Investigating what happens when aqueous solutions are electrolysed using inert electrodes 4.17 Key concept: Electron transfer, oxidation and reduction 4.18 Maths skills: Make order of magnitude calculations 4.2
132 134 136 138 140 142 144
146 148
150 152 154 156 158 160
162 164 166
Chapter 5 Energy changes 172
Key concept: Endothermic and exothermic reactions 174 5.2 Required practical: Investigate the variables that affect temperature changes in reacting solutions such as, acid plus metals, acid plus carbonates, neutralisations, displacement of metals 176 5.1
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5.3 5.4 5.5 5.6 5.7
Reaction profiles Energy change of reactions Cells and batteries Fuel cells Maths skills: Recognise and use expressions in decimal form
178 180 182 184 186
Chapter 6 The rate and extent of chemical change 192
6.1 6.2 6.3 6.4 6.5
6.6 6.7 6.8 6.9 6.10 6.11 6.12 6.13 6.14
Measuring rates 194 Key concept: Limiting reactants and molar masses 196 Calculating rates 198 Factors affecting rates 200 Required practical: Investigate how changes in concentration affect the rates of reactions by a method involving the production of gas and a method involving a colour change 202 Factors increasing the rate 204 Collision theory 206 Catalysts 208 Reversible reactions and energy changes 210 Equilibrium 212 Changing concentration and equilibrium 214 Changing temperature and equilibrium 216 Changing pressure and equilibrium 218 Maths skills: Use the slope of a tangent as a measure of rate of change 220
Chapter 7 Hydrocarbons
226
228
7.1
Crude oil, hydrocarbons and alkanes Fractional distillation and petrochemicals 7.3 Properties of hydrocarbons 7.4 Combustion 7.5 Cracking and alkenes 7.6 Structure and formulae of alkenes 7.7 Reactions of alkenes 7.8 Alcohols 7.9 Carboxylic acids 7.10 Addition polymerisation 7.11 Condensation polymerisation 7.12 Amino acids 7.13 DNA and other naturally occurring polymers 7.14 Key concept: Intermolecular forces 7.15 Maths skills: Visualise and represent 3D models
252 254 256
Chapter 8 Chemical analysis
262
264 266 268
Key concept: Pure substances 8.2 Formulations 8.3 Chromatography 8.1
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270 272 274 276 278
280 282 284 286
Chapter 9 The atmosphere
292
294 296 298 300 302 304 306 308
7.2
230 232 234 236 238 240 242 244 246 248 250
Required practical: Investigate how paper chromatography can be used in forensic science to identify an ink mixture used in a forgery 8.5 Test for gases 8.6 Flame tests 8.7 Metal hydroxides 8.8 Tests for anions 8.9 Required practical: Use chemical tests to identify the ions in unknown single ionic compounds 8.10 Instrumental methods 8.11 Flame emission spectroscopy 8.12 Maths skills: Use an appropriate number of significant figures 8.4
9.1
Proportions of gases in the atmosphere The Earth’s early atmosphere 9.3 How oxygen increased 9.4 How carbon dioxide decreased 9.5 Key concept: Greenhouse gases 9.6 Human activities 9.7 Global climate change 9.8 Carbon footprint and its reduction 9.9 Limitations on carbon footprint reduction 9.10 Atmospheric pollutants from fuels 9.11 Properties and effects of atmospheric pollutants 9.12 Maths skills: Use ratios, fractions and percentages 9.2
310 312 314 316
Chapter 10 Sustainable development 322
Key concept: Using the Earth’s resources and sustainable development 324 10.2 Potable water 326 10.3 Required practical: Analysis and purification of water samples from different sources, including pH, dissolved solids and distillation 328 10.4 Waste water treatment 330 10.5 Alternative methods of metal extraction 332 10.6 Life Cycle Assessment and recycling 334 10.7 Ways of reducing the use of resources 336 10.8 Corrosion and its prevention 338 10.9 Alloys as useful materials 340 10.10 Ceramics, polymers and composites 342 10.11 Haber process 344 10.12 Production and use of NPK fertilisers 346 Maths skills: Translate information 10.13 between graphical and numerical form 348 10.1
Glossary 355 Index 361
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Chemistry
ATOMIC STRUCTURE AND THE PERIODIC TABLE ideas you HaVe met Before:
eLements, mixtures and ComPounds • Mixtures can be separated easily by fi ltering and other ways. • Elements cannot be broken down by chemical means. • Compounds are made from elements chemically combined.
atoms and tHeir struCture • Electrons have a negative charge. • Atoms have a nucleus with a positive charge. • Electrons orbit the nucleus in shells.
C
some eLements and tHeir ComPounds • Helium is unreactive and used in balloons. • Sodium chloride is used to fl avour and preserve food. • Chlorine is used to kill bacteria in swimming pools.
metaLs in tHe PeriodiC taBLe • Gold, silver and platinum are precious metals. • Mercury is a liquid metal. • Zinc, copper and iron are used to make many useful objects.
metaLs and non-metaLs • Gold, iron, copper and lead are metals known for centuries. • Oxygen and nitrogen are gases of the air. • Sulfur is a yellow non-metal.
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1 in tHis CHaPter you wiLL find out aBout:
wHat modeL do we use to rePresent an atom? • Electrons fi ll the shells around the nucleus in set pattern orders. • Protons and neutrons make up the nucleus. • Electrons can be lost from or gained into the outer shell.
Li
How did tHe modeL of tHe atom deVeLoP? • Atoms used to be thought of as small unbreakable spheres. • Experiments led to ideas of atoms with a nucleus and electrons. • Electrons in shells and the discovery of the neutron came later.
wHy Can we use CarBon datinG? • Atoms of an element always have the same number of protons. • They do not always have the same numbers of neutrons. • Elements exist as different isotopes.
C
C
Spot the difference in these isotopes
wHy is HeLium so unreaCtiVe and sodium so reaCtiVe? • The outer shell of helium can take no more electrons. • The outer shell of sodium has 1 electron which it needs to lose. • Metals need to lose electrons, non-metals do not.
wHat do transition metaL ComPound soLutions LooK LiKe? • Transition metals are harder and stronger than Group 1 metals. • Transition metals are often used as catalysts. • Transition metal compounds often form coloured solutions.
Atomic structure and the periodic table
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Chemistry
Elements and compounds
Key words
Learning objectives:
compound element
• identify symbols of elements from the periodic table • recognise compounds from their formula • identify the elements in a compound.
All the elements are listed in the periodic table. The elements in the formulae of any compound, no matter how large, can be identified by using the periodic table.
Elements and compounds An element is a substance that cannot be broken down chemically. A compound is a substance that contains at least two different elements, chemically combined in fixed proportions.
1 1
PERIODIC TABLE ELEMENTS 1-20
H
hydrogen
7 3
Li
lithium
23 11
9 4
Be
beryllium
24
Na 12 Mg
sodium magnesium
39 19
K
40 20
11 5
B
boron
27 13
Al
12 6
C
14 7
N
16 8
O
4 2
19 9
F
carbon
nitrogen
oxygen
fluorine
28 14
31 15
32 16
35 17
Si
P
aluminium silicon phosphorus
S
sulfur
He
helium
Cl
chlorine
20 10
Ne
neon
40 18
Ar
argon
Ca
potassium calcium
Figure 1.1 Two sections of the periodic table. Can you find magnesium and oxygen?
+ Figure 1.2 Magnesium (metal) reacts with oxygen (gas) to make magnesium oxide (compound).
+ x z y
Figure 1.3 The reaction of the elements in Figure 1.2 can be represented by models of their atoms. 1
Identify the following substances as elements or compounds. copper copper chloride copper sulfate
2
14
Name the elements in beryllium chloride.
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Compounds and elements
1.1
Copper is an element. It cannot be broken down into any other substances, but salt can. The chemical name for common salt is sodium chloride. Sodium chloride is a compound. Sodium chloride can be broken down to make sodium and chlorine, but this is not easy to do because the sodium and chlorine are chemically combined. We need to use electricity to make sodium and chlorine from sodium chloride. Sodium and chlorine cannot be broken down any further. Sodium and chlorine are elements. There are only about 100 elements but these can join together chemically to make an enormous number of compounds. They need chemical reactions to do this. 3
Identify the elements in potassium bromide.
4
Predict the products when lead iodide is split by electricity.
Key information When chlorine reacts to make a compound it chemically combines and becomes a chloride. Similarly, bromine reacts to become a bromide and oxygen reacts to become an oxide.
Making the element copper into a compound A chemical reaction is needed to make copper (an element) into a compound. If copper is burned in oxygen it forms the compound copper oxide. Chemical reactions always involve the formation of one or more new substances, and often involve a detectable energy change.
Did you know? Compounds can only be separated into elements by chemical reactions. To get the element copper back, the oxygen needs to be chemically removed. This is done using hydrogen. The oxygen combines with the hydrogen to make water, another compound. copper oxide + hydrogen → copper + water Figure 1.4 Copper (an element) burning in oxygen (an element) to make copper oxide (a compound). This is normally done by heating the copper in crucibles. 5
What is the name of the compound made from sodium and oxygen.
6
Oxygen can be removed from iron(III) oxide by carbon monoxide. Identify the element and compound produced.
7
Substance D reacted with hydrogen to form zinc and water. Explain whether substance D is an element or compound. Google search: 'unusual periodic table (images)'
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Chemistry
Atoms, formulae and equations Learning objectives: • explain that an element consists of the same type of atoms • explain that atoms join together to make molecules • explain how formulae represent elements and compounds.
Key words balanced compound element equation
formula molecule symbol
All substances are chemicals. Many people say “I don’t want chemicals in my food” not realising that foods are chemicals. Our food is made of compounds and mixtures. Compounds are elements joined together in many different ways. In fact, we are all made of chemical compounds.
Atoms and molecules Elements are made up of atoms that are all the same. Compounds are made from atoms (or charged atoms) of different elements which have been chemically joined together. CH3COOH is a compound made from three elements. These are carbon, C, hydrogen, H, and oxygen, O. You will know it as vinegar. Its chemical name is ethanoic acid. The atoms join by sharing their electrons.
Figure 1.5 Vinegar is used in salad dressing. Its chemical name is ethanoic acid. Ethanoic acid is a compound made from carbon, hydrogen and oxygen.
ethanoic acid
Figure 1.7 The molecule in vinegar CH3COOH oxygen molecule
Figure 1.6 Gold is an element. Carbon monoxide is a compound. How could you tell by looking at these diagrams?
If two or more atoms join together by sharing their electrons the atoms form a molecule. Two examples of molecules are oxygen and water. 1
Explain whether the following substances are elements or compounds. C (carbon), CO2 (carbon dioxide), Cl2 (chlorine) and SO3 (sulfur trioxide)
2
For each of the substances listed below, give: a the names of the elements b how many different types of atoms they contain c how many atoms are in the molecule overall.
oxygen molecule
water molecule
Figure 1.8 Which of these is a molecule of an element? How can you tell? water molecule
Key information In a chemical formula, the subscript number on the bottom right is the number of atoms in the molecule, for example C2H6 has two carbon atoms joined to six hydrogen atoms.
CO2 S8 Cl2 SO3 C60 C 4H10
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Formulae
1.2
You can see which elements are in a compound by looking at its formula. For example, the compound magnesium oxide (MgO) contains Mg (magnesium) and O2 (oxygen). Look back at the section of the periodic table on page 14. Elements from Group 1 and elements from Group 7 combine to make compounds in a fixed ratio of 1 : 1. • The formula of lithium fluoride is LiF. Elements from Group 2 and elements from Group 6 also combine to make compounds in a fixed ratio of 1 : 1.
Did you know?
• The formula of calcium oxide is CaO.
Oxygen exists as a pair of atoms not as a single atom. Its formula is O2. In a symbol equation, O2 must be written not just O.
Elements from Group 2 and elements from Group 7 combine to make compounds in a fixed ratio of 1 : 2. The element that you need two of has a suffix ‘2’ after the symbol. • The formula of calcium fluoride is CaF2. Elements from Group 1 and elements from Group 6 combine to make compounds in a fixed ratio of 2 : 1. Again, the element that you need two of has a suffix ‘2’ after the symbol. • The formula of sodium sulfide is Na2S. 3
Give the names of the elements in MgSO4 and in CH4.
4
Determine the formulae of lithium chloride, magnesium chloride and potassium oxide.
Equations and balancing When magnesium reacts with oxygen it makes magnesium oxide. We can represent this chemical change using word equations and balanced symbol equations.
Key information
The equations show us what substances we start with (reactants) and what new substances they turn into (products).
In a balanced equation ‘O2’ means a pair of atoms joined together in a molecule. ‘2Mg’ means two separate atoms, where the ‘2’ is added to balance the equation.
The word equation is: magnesium +
oxygen → magnesium oxide
The symbol equation is:
Mg
+
O2
→
MgO
1
2
1 1
✗
This is not a balanced symbol equation because there are two atoms of oxygen in the reactants and only one atom of oxygen in the products. We need to add ‘2’ to the front of the formulae of magnesium and magnesium oxide. This gives:
2Mg
+
O2
→
2MgO
2
2
2 2
✓
7
5
Write a balanced equation for the formation of sodium chloride from sodium, Na and chlorine, Cl2.
6
Write a balanced equation for the formation of aluminium oxide Al2O3.
Complete and balance the following equation by suggesting values for D, E and F: D N2 + E O2 → F NO2
Google search: 'unusual periodic table (images)'
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Chemistry
Mixtures Learning objectives: • recognise that all substances are chemicals • understand that all substances are either mixtures, compounds or elements • explain that mixtures can be separated.
Key words chromatography filtration mixture separation
You will have begun to use of a range of equipment to safely separate chemical mixtures and we need to extend this range of techniques. Filtering and distillation are probably familiar but fractional distillation can also be used to separate mixtures.
Mixtures Many substances are made of mixtures. Mixtures can easily be separated because the chemicals in them are not joined together. Mixtures can be separated by filtration, crystallisation, simple distillation, fractional distillation and chromatography.
Figure 1.9 Separating mixtures
Let’s take a mixture of salt and copper. To separate them we add water to the mixture. The salt dissolves but the copper does not. The salt solution can be filtered through a filter paper, leaving the copper behind as a residue. The salt solution can be crystallised to make solid salt crystals. These physical processes do not involve chemical reactions and no new substances are made. After separating, salt and copper can be mixed again.
18
1
Draw a diagram of the equipment used to filter salt solution from copper.
2
Explain why a blend of copper and salt is not a compound.
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Separating mixtures A mixture consists of two or more elements or compounds not chemically combined together. The chemical properties of each substance in the mixture are unchanged. Mixtures can be separated by physical processes. These processes do not involve chemical reactions. These separation processes include: • fi ltration • crystallisation filtration
• distillation • chromatography
crystallisation
distillation
chromatography
crystallising dish
1.3 Key information filtration separates insoluble substances from soluble substances and distillation separates liquids that have different boiling points.
crystals
Figure 1.10 Techniques for separating mixtures. 3
Which technique would be used to separate coloured inks in a mixture?
4
Which technique would be used to separate alcohol (boiling point 80°C) and water?
Fractional distillation Some mixtures are very complex and have many different components. These mixtures can be separated either (a) by using different techniques in sequence or (b) by the same technique which includes multiple separations. An example of the fi rst approach is fi ltration followed by crystallisation.
did you Know? fractional distillation is used to separate the different substances in crude oil.
fractional distillation
An example of the second approach is fractional distillation, where different liquids with different boiling points are separated at different points in the process. Fractional distillation works by using a tall tower of gaps and surfaces, which are gradually colder towards the top. The liquid mixture is heated at the bottom and the liquids boil together to make a mixture of gases. As each gas reaches a surface at the same temperature as its boiling point (or condensing point) that gas will condense and the condensed liquid will run off. The other gases continue up through the gaps until they reach the surface at their condensing temperature. Eventually nearly all the gases in the mixture will condense and be collected as separated liquids. The fi nal gas is left at the top of the tower and is collected as a gas. 5
Suggest how you would collect a specimen of clean copper sulfate crystals from a mixture of solid copper sulfate, sand and alcohol.
6
Suggest the order (from bottom to top) in which these liquids would be collected from the fractional distillation tower: a (boiling point 85 °C)
c
(boiling point 35 °C)
b (boiling point 100 °C)
d (boiling point 165 °C)
Figure 1.11 Fractional distillation of mixtures with different boiling points
Google search: ‘chromatography, filtration, fractional distillation’
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Chemistry
Changing ideas about atoms Learning objectives: • describe how the atomic model has changed over time • explain why the atomic model has changed over time • understand that a theory is provisional until the next piece of evidence is available.
Key words Ernest Rutherford Geiger and Marsden J. J. Thomson James Chadwick John Dalton Niels Bohr
The idea of atoms has changed hugely over the years. At the moment, scientists believe atoms are very small, have a very small mass and are made of protons, electrons and neutrons. Our current theories were developed by imagination, evidence and advances in technology, with each new idea being built on the ideas of earlier scientists.
Developing the atomic theory Explanations about atoms began about 400 BC, when the Greek philosopher Democritus described materials as being made of small particles which could not be divided. He called these particles ‘atoms’. However, he had no evidence. It was just an idea. Little more was suggested for more than 2000 years, but in 1803 the British scientist John Dalton used his observations to describe the atom in more detail. His model described an atom as a ‘billiard ball’ which could not be divided.
Figure 1.12 Dalton’s idea of atoms: they were like tiny billiard balls.
Dalton’s model was then changed as new evidence was found. In 1897, 94 years later, J. J. Thomson discovered the electron. Thomson developed the way that the atom was thought of by using a ‘plum pudding’ model to describe atoms. Negative electrons were thought to be embedded in a ball of positive charge, rather like the fruit (the electrons) are part of a pudding (the ball of positive charge).
20
1
Suggest why Dalton’s atomic model did not include positive and negative charge.
2
Explain why the discovery of the electron changed the Dalton model of the atom.
Key information At each stage, the explanations of atomic theory were provisional until more convincing evidence was found to make the model better.
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Changing theories
1.4
Sometimes ideas can develop rapidly because of unexpected results. In 1909 Geiger and Marsden had really surprising results in their experiment with gold leaf and alpha particles. These results led Geiger, Marsden and Ernest Rutherford to propose a new idea that an atom has a nucleus. In 1911, Rutherford suggested the atom had a positively charged nucleus in the centre where most of the mass of the atom was concentrated and much of the atom was empty space. This was the nuclear model of the atom.
Figure 1.13 Rutherford and Geiger in their lab in Manchester, UK.
In 1913, Niels Bohr used theoretical calculations that agreed with experimental evidence to adapt the nuclear model. He explained that the electrons orbited the nucleus in definite orbits at specific distances from the nucleus. He explained that a fixed amount of energy (a quantum of energy) is needed for an electron to move from one orbit to the next. Electrons only exist in these orbits. 3
Suggest why Bohr proposed that electrons orbited the nucleus in shells.
4
What is meant by the phrase ‘quantum of energy’?
Further development of atomic theory Later experiments gradually led to the idea that the positive charge of any nucleus can be sub-divided into a whole number of smaller particles. Each of these particles had the same amount of positive charge. In 1920 the term ‘proton’ was first used in print for these particles. In 1932, James Chadwick discovered the neutron. Again this discovery involved experimental evidence and mathematical analysis. 5
Draw a timeline of the discoveries that led to our present understanding of the atomic theory.
6
Suggest why it was twelve years between finding protons and neutrons.
Did you know? As a challenge you can find out about the Geiger and Marsden experiment that changed the theory from a ‘plum-pudding’ atom to a nuclear atom, it is a famous turning point in the understanding of atoms.
Did you know? The idea of atoms as small particles is not new. However, our ideas about the theory of atoms are still developing. Search on ‘CERN LHC’ to find out more.
Google search: 'timeline of development of atomic theory'
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Chemistry
Modelling the atom Learning objectives: • describe the atom as a positively charged nucleus surrounded by negatively charged electrons • explain that most of the mass of an atom is in the nucleus • explain that the nuclear radius is much smaller than that of the atom and with most of the mass in the nucleus.
Key words charge electron electron shell negative nucleus positive
Atoms are the building blocks of all matter, both living and non-living, simple and complex. Atoms join together in millions of different ways to make all the materials around us. We can explain how everything, including ourselves, is made by using ideas and models of atoms.
Atoms Individual atoms are very small. There are about ten million million atoms in this full stop. An atom is made up of a nucleus that is surrounded by electrons. • The nucleus carries the positive charge. • Electrons, which surround the nucleus, each carry a negative charge.
Figure 1.14 Image of gold atoms. Magnification x 16 000 000.
It depends how it is measured but the diameter of an atom is about 10 –10 m. That’s 0.000 000 1 mm. If we imagine that an atom is blown up to the size of a football stadium the nucleus would be the size of a peanut placed on the centre spot. 1
What is the type of charge in the nucleus?
2
Helium has two positive charges in the nucleus. Predict the number of electrons in a helium atom.
Figure 1.15 The structure of a hydrogen atom. What charge does an electron carry?
More on atoms Electrons occupy the space around the nucleus in ‘shells’. The space between the nucleus and the electron shells is empty space.
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The nucleus contains most of the mass of the atom and the electrons contribute very little. On the other hand, the radius of the atom, where the electrons are orbiting, is much larger than the radius of the nucleus in the centre.
1.5
When we are talking about these differences we are talking about small sizes. Atoms are very small. A typical atomic radius is about 0.1 nm (1 × 10 −10 m). The radius of a nucleus is less than one ten-thousandth of the radius of an atom (about 1 × 10 −14 m). Typical atomic radius 1 × 10
−10
m
Typical radius of a nucleus 1 × 10 −14 m
The radius of an atom is measured in many different ways. This is because the outer electron shell is not a fixed boundary, and so its position can only be measured approximately. 3
Most of the atom is empty space. What does this suggest about the size of an electron?
4
Explain why the radius of the nucleus is much smaller than the radius of the whole atom.
Did you know? An atom of gold has a mass of about 3.3 × 10 −22 g and a radius of about 1.4 × 10 −10 m. Most of the mass of the atom is in the middle, in the nucleus.
Atomic radii Atoms are very small. A typical atomic radius is about 0.1 nm (1 × 10 −10 m). However, the radius of an atom increases within a group of elements. For example the atomic radii of Li, Na, K, increase as more electrons are ‘added’ to the atom. 5
Suggest why the radius of potassium is larger than the radius of lithium.
6
The positive charge on a Li nucleus is 3. The positive charge on a Ne nucleus is 10. As more negative electrons are added one by one to atoms from Li up to Ne the radius gets smaller, not bigger. Suggest why. Use ideas about opposite charges.
Key information Remember that the typical radius of a nucleus is less than 1/10 000 th of the typical radius of an atom.
Google search: 'atomic radius Royal Society of Chemistry'
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Chemistry
Relating charges and masses Learning objectives: • describe the structure of atoms • recall the relative masses and charges of protons, neutrons and electrons • explain why atoms are neutral.
We have seen how ideas about atoms have changed over the years. Currently, scientists believe atoms are made of three important particles – protons, electrons and neutrons. • The number of protons and neutrons are important in nuclear reactions. • The number of electrons is important in chemical reactions.
Structure of atoms An atom is made up of a nucleus that is surrounded by electrons.
Key words atomic number electron neutral neutron proton
Did you know? Even these protons and neutrons can be broken down further in huge particle accelerators such as the one built deep underneath Switzerland by a joint team of scientists and engineers from many European countries.
• The nucleus carries a positive charge. • The electrons that surround the nucleus each carry a negative charge. The nucleus of an atom is made up of protons and neutrons. • Protons have a positive charge. • Neutrons have no charge. An atom always has the same number of protons (+) as electrons (–) so atoms are always neutral. The atomic number is the number of protons in an atom. The atomic number for helium is 2 because it has two protons. 1
Lithium has an atomic number of 3. Predict the number of electrons in lithium.
2
The neon atom has 10 protons. Explain why the neon atom is neutral.
3
Use the periodic table to identify the element with 3 protons.
4
Determine the number of protons in an atom of calcium, Ca.
Key information It is because a helium atom has two protons that it has an atomic number of 2.
Masses and charges The nucleus of an atom is made up of particles (protons and neutrons) that are much heavier than electrons. The relative masses and charges of electrons, protons and neutrons are shown in the table.
24
Figure 1.16 The structure of a helium atom. There are the same number of protons and electrons.
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Relative charge –1
0.0005
Proton
+1
1
0
1
Neutron
1.6
Relative mass
Electron
Did you know?
5
A fluorine atom has 9 positive charges, 9 negative charges and a mass of 19. Describe the structure of its atom.
6
A chlorine atom has 17 electrons and a mass of 35. Describe the structure of its atom.
Electrons have such a small relative mass that it is usually treated as zero.
Losing electrons If an atom has an atomic number of 3 and a neutral charge, it must be a lithium atom. It has a neutral charge because the atom has three protons (+) and three electrons (–). If the lithium atom loses one negatively charged electron it then becomes a charged particle with one positive charge that is not balanced out by a negative charge. Atomic number
Number of protons
Number of electrons
Lithium atom
3
3
3
0
Lithium charged particle
3
3
2
+1
Figure 1.17 A neutral lithium atom loses an electron and becomes charged.
Charge
If an atom loses electrons and becomes charged, this charged particle is called a positive ion. 7
If a magnesium atom loses 2 electrons it becomes a charged particle. It still has a mass of 24. Write out the atomic number, number of protons, number of electrons, number of neutrons and the charge of a a magnesium atom b a magnesium ion.
8
Explain why a magnesium atom is neutral but a magnesium ion is charged.
9
Nitride ions have a –3 charge. Work out the number of electrons in a nitride ion, given that the atomic number of nitrogen is 7.
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Chemistry
Sub-atomic particles
Key words
Learning objectives: • use the definitions of atomic number and mass number • calculate the number of protons, neutrons and electrons in atoms • calculate the number of sub-atomic particles in isotopes and ions.
atomic mass isotope neutrons protons
Smoke detectors, archaeological dating and bone imaging all use isotopes. Some elements have more than one type of atom. These different types of atom have different numbers of neutrons and are called isotopes.
Atomic number and mass number The nucleus of an atom is made up of protons and neutrons. • The atomic number is the number of protons in an atom. • The mass number of an atom is the total number of protons and neutrons in an atom. If a particle has an atomic number of 11, a mass number of 23 and a neutral charge, it must have: • 11 protons, because it has an atomic number of 11. • 11 electrons, because there are 11 protons and the atom is neutral. • 12 neutrons, because the mass number is 23 and there are already 11 protons (23 – 11 = 12). Here are some more examples. Atomic number
Mass number
Number of protons
Number of electrons
Number of neutrons
Carbon
6
12
6
6
6
Fluorine
9
19
9
9
10
Sodium
11
23
11
11
12
Aluminium
26
1
Complete the row for an atom of aluminium, Al.
2
Work out the number of protons, electrons and neutrons in an atom with an atomic number of 15 and a mass number of 31.
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Isotopes
1.7
All atoms of carbon have 6 protons, so its atomic number is 6. Most carbon atoms have 6 neutrons, so the mass number is 6 + 6 = 12. This form of the carbon atom is written as 126 C. Another form of carbon, 146 C, has an atomic number of 6 (6 protons) and a mass number of 14. It must therefore have 8 neutrons (14 – 6). 146 C is sometimes written as carbon-14. 12 C and 146 C are isotopes of carbon. 6 3
Write the isotope symbol for an atom that has 17 protons and 18 neutrons.
4
Identify all the sub-atomic particles in an atom of carbon-13.
Key information In the symbol 126C, the smaller number (6) is the atomic number and the larger number (12) is the mass number.
Relative abundance of isotopes Most elements have two or more isotopes. For example, hydrogen has three common isotopes. Isotope
Electrons
Protons
Neutrons
Mass number
H
1
1
0
1
H
1
1
1
2
H
1
1
2
3
1 1 2 1 3 1
Did you know?
The relative atomic mass of an element is the average mass of the different isotopes of an element. Chlorine’s Ar of 35.5 is an average of the masses of the different isotopes of chlorine. Cl and 37 Cl. If there were 50% There are two main isotopes 35 17 17 of each of the isotopes what would be the average mass? The Cl isotopes. So we need answer is 36. But there are less of the 37 17 a relative abundance calculation:
Ar =
mass of first isotope × mass of second isotope × % of first isottope + % of second isotope
The mass numbers of the isotopes of hydrogen are 1, 2 and 3. However, there are not equal proportions of each type of isotope in a sample of hydrogen gas, so the average atomic mass of hydrogen is 1.008.
100
For example, for chlorine: the abundance values are: 75% 35 Cl and 25% 37 Cl 17 17 Therefore: Ar = (75 × 35) + (25 × 37) 100 2625 + 925 = 100 = 35.5 5
Explain the similarities and differences between the three isotopes of hydrogen.
6
Element X has two isotopes, mass 27 and 29. Calculate the relative atomic mass of X if the first isotope has abundance of 65% and the second isotope has 35% abundance.
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Chemistry
Electronic structure
Key words
Learning objectives: • explain how electrons occupy ‘shells’ in an order. • describe the pattern of the electrons in shells for the first 20 elements.
electronic structure electron shells energy levels group
The electrons of an atom are arranged in patterns. The electrons fill up shells in order until that shell can take no more electrons. The next electron goes into the next shell. These patterns are the key to the behaviour of atoms.
The ‘build-up’ of electrons Electrons occupy shells around the nucleus. The electron shell nearest to the nucleus takes up to two electrons. The second shell takes up to eight electrons. The next electrons occupy a third shell. Oxygen has an atomic number of 8. It has eight protons and so it has eight electrons in the space around the nucleus. The first two go into the first shell. As the first shell is now full, the next 6 electrons go into the second shell. The electron pattern for oxygen is then 2,6. 1
Draw the electron pattern for hydrogen and for lithium.
2
Write down the electron pattern for nitrogen.
The shells are not fixed rings and are also known as energy levels.
Figure 1.18 For oxygen, the eight electrons can written as 2,6 or be drawn like this.
Electron patterns and groups The periodic table is arranged in order of ‘proton number’. There is a very important link between electronic structure and the periodic table. For example, let us consider an atom of an element that has the electronic structure of 2,8,6. • This element has three electron shells, so it is in the third row of the periodic table. • It has six electrons in its outer shell, so it is in the sixth column. • Using the periodic table, we can find that its atomic number is 16 and it is the element sulfur, S. H
He
hydrogen
helium
1
Li
Be
lithium
2
B
C
N
O
F
Ne neon
beryllium
boron
carbon
nitrogen
oxygen
fluorine
3
4
5
6
7
8
9
10
Na
Mg
Al
Si
P
S
Cl
Ar
sodium
magnesium
aluminium
silicon
phosphorus
sulfur
chlorine
argon
11
12
13
14
15
16
17
18
K
Ca
potassium
calcium
19
20
this row has the elements of period 3
Figure 1.19 Period 3 contains the elements from sodium to argon.
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We can also work the other way. Find the element with the atomic number 12 in the periodic table. This is magnesium, Mg.
1.8
• It is in the third row, so it has three electron shells. • It is in the second column, so it has two electrons in its outer shell. • It has the electronic structure of 2,8,2. The column of elements is known as a group. So column 2 is Group 2. 3
An atom of an element has an atomic number of 11. Figure 1.20 An atom of fluorine, 2,7.
a Draw a diagram to show the pattern of electrons. b Identify the element. c Identify the group to which it belongs. 4
Work out the electronic structure of the element that has an atomic number 9.
5
An element has a mass number of 40 and an electron arrangement of 2,8,8,2. Identify the element and work out the number of neutrons.
Maximum numbers The electronic structure of each of the first 20 elements can be worked out using: • the atomic number of the element • the maximum number of electrons in each shell. The third shell takes up to eight electrons before the fourth shell starts to fill. Element 19, potassium, has the electronic structure 2,8,8,1.
Did you know? It took many years for scientists to work out the theory of electrons occupying shells. They started with the behaviour of elements and then looked for patterns.
Use the periodic table to help you to answer these questions. 6
Work out the electronic structure of argon.
7
Use a blank periodic table sheet to draw out the electronic structure of the first 20 elements, putting the diagram in the correct box. What do you notice about the group number and the number of electrons in the outer shell?
8
Key information
The electronic configuration for the ion of an unknown isotope 26 3+ X is 2,8. a Work out the atomic number of element X. b Determine the number of neutrons in X.
Figure 1.21 An atom of phosphorus, 2,8,5.
Do not try to use this method to work out the electronic structure of gold, as it has 79 electrons.
c Explain which group in the periodic table element X is in.
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Chemistry
The periodic table
Key words
Learning objectives:
electron shells energy levels group period
• explain how the electronic structure of atoms follows a pattern • recognise that the number of electrons in an element’s outer shell corresponds to the element’s group number • explain that the electronic structure of transition metals positions the elements into the transition metal block.
We know that the periodic table is arranged in rows and columns and the elements are written in order of their atomic number. So why are all the elements in the last column all unreactive gases? Why are all the elements in the first column highly reactive metals? The answer lies in the pattern of their electrons.
The order of elements and electron patterns As we have seen, the elements in the periodic table are arranged in order of atomic number. Atomic number is the number of protons in an atom. As atoms are neutral, the atomic number also gives the number of electrons in an atom.
Key information Remember: • the first shell of electrons carries up to 2 electrons • the second shell carries up to 8 electrons. • the third shell carries up to 8 electrons and the fourth up to 18.
For example, the atomic number of hydrogen is 1, carbon is 6 and sodium is 11. This means that hydrogen is the first element in the table, carbon is the sixth and sodium is the eleventh. We have also seen that electrons occupy energy levels (or shells). Each element has a pattern of electrons (known as its electronic structure) that is built up in a particular order. The electronic structure of each of the first 20 elements can be worked out using: • the atomic number of the element • the maximum number of electrons allowed in each shell.
Figure 1.22 The electron pattern of a lithium atom
The third shell takes up to eight electrons before the fourth shell starts to fill. Element 20, therefore, has the electronic structure 2,8,8,2. 1
Which element has the atomic number 13?
2
Which element has an electronic structure of 2,8,7?
Arrangement of groups and periods A row of the periodic table is known as a period. The first row of the periodic table contains the elements hydrogen and helium. This row is called the first period. These two elements only have electrons in the first shell (energy level). Lithium’s third electron goes into the next electron shell. Lithium starts a new row in the periodic table. This second row is called the second period.
30
the atomic number of hydrogen is 1
Li Be 3
4
H
He
1
2
B C N O F Ne 5
6
7
8
9
10
Na Mg Al Si P S Cl Ar 11
12
13
14
15
16
17
18
K Ca 19
20
Figure 1.23 Lithium has an atomic number of 3. It has three protons and three electrons.
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Let’s consider an atom of the element that has the electronic structure of 2,8,2. We can work out that it is in the third row of the periodic table. This atom has three electron shells. It has two electrons in its outer shell. Its atomic number is 12. (You can work this out by adding up the number of electrons.) Looking at the periodic table, the atom is of the element magnesium, 12. So the element is therefore, in the third period. Looking again, this atom only has two electrons in its outer shell, so it is in the second column. A column is known as a group. This second column is known as Group 2. The group number refers to the number of electrons in the outside shell. We have seen that the electron pattern for Li is 2,1. We can work out that the pattern for Na (11 electrons) is 2,8,1 and K (19 electrons) is 2,8,8,1. Both of these elements have one electron in their outside shell. They are both in the first column. This column is known as Group 1. 3 4
1.9 Key information The periodic table is arranged in rows (called periods). Each period number is the same as the number of electron shells. Each period contains the elements whose outside shell of electrons is ‘filling up’.
What is the pattern of electrons in the atom of the element with an atomic number 16? Identify the element, its group and its period.
this column has the elements of group 1
H
hydrogen 1
To which group and period does the element chlorine belong? Li
Electronic structure and behaviour of elements Element 20 has the electronic structure 2,8,8,2 and is in Group 2. The fourth shell can take up to 18 electrons, but the next element (element number 21) is not in Group 3. Instead the next 10 elements (numbers 21 to 30) in period 4 are part of the first transition element period. Transition elements have characteristics that are similar to each other. For example, even though they are all metals, iron behaves more like copper than sodium.
lithium 3
Na
sodium 11
K
potassium 19
Be
beryllium 4
Mg
magnesium 12
B
boron 5
Al
aluminium 13
C
carbon 6
Si
silicon 14
Ca
calcium 20
Figure 1.24 These elements are in the first column as they all have one electron in their outside shell.
Elements in groups often behave similarly to other elements in that group: • the elements in Group 1 are all highly reactive metals • the elements in Group 7 (the halogens) all react with Group 1 metals to make salts • the elements in Group 0 are all unreactive (or noble) gases. This pattern in reactivity occurs because the elements in each group have the same number of electrons in their outer shells. Use the periodic table to help you to answer these questions. 5
What are the electronic structures of F and Cl? Why are they both found in Group 7?
6
Identify the elements in Group 6.
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