Cambridge international AS a level chemistry coursebook 3e by lawrie ryan, roger norris

In this chapter you will learn how to: describe the structure of the atom as mostly empty space surrounding a very small nucleus that consists of protonsand neutrons and state that elect

Chemistry for Cambridge International AS & A Level

COURSEBOOK Lawrie Ryan & Roger Norris

Teachers play an important part in shaping futures Our Dedicated Teacher Awards recognise the hard work that

teachers put in every day.

Thank you to everyone who nominated this year, we have been inspired and moved by all of your stories Well done to all of our nominees for your dedication to learning and for inspiring the next generation of thinkers,

leaders and innovators.

For more information about our dedicated teachers and their stories, go to

dedicatedteacher.cambridge.org

Throughout this book, you will notice lots of different features that will help your learning These areexplained below

LEARNING INTENTIONS

These set the scene for each chapter, help with navigation through the coursebook and indicate theimportant concepts in each topic

BEFORE YOU START

This feature contains questions and activities on subject knowledge you will need before starting thischapter

SCIENCE IN CONTEXT

This feature presents real-world examples and applications of the content in a chapter, encouragingyou to look further into topics There are discussion questions at the end which look at some of thebenefits and problems of these applications

PRACTICAL ACTIVITIES

This book does not contain detailed instructions for doing particular experiments, but you will findbackground information about the practical work you need to do in these boxes There are also twochapters, P1 and P2, which provide detailed information about the practical skills you need to developduring the course

Questions

Appearing throughout the text, questions give you a chance to check that you have understood the topicyou have just read about You can find the answers to these questions in the digital version of the

coursebook

KEY DEFINITIONS

Key definitions for important scientific principles, laws and theories are given in the margin and

highlighted in the text when it is first introduced You will also find these definitions in the Glossary atthe back of this book

KEY WORDS

Key vocabulary is highlighted in the text when it is first introduced Definitions are then given in themargin, which explain the meanings of these words and phrases

You will also find definitions of these words in the Glossary at the back of this book

COMMAND WORDS

style questions when they are first introduced In the margin, you will find the Cambridge Internationaldefinition*

Command words that appear in the syllabus and might be used in exams are highlighted in the exam-You will also find the same definitions in the Glossary at the back of this book

There is a summary of key points at the end of each chapter

SELF-EVALUATION

The summary checklists are followed by ‘I can’ statements which match the learning intentions at thebeginning of the chapter You might find it helpful to rate how confident you are for each of these

www.cambridgeinternational.org

The resource index is a convenient place for you to download all answer files for this resource

We hope not only that this book will help you to succeed in your future studies and career, but also that itwill stimulate your curiosity and your imagination Today’s students become the next generation of

scientists and engineers, and we hope that you will learn from the past to take chemistry to ever-greaterheights

In this chapter you will learn how to:

describe the structure of the atom as mostly empty space surrounding a very small nucleus that consists of protonsand neutrons and state that electrons are found in shells in the space around the nucleus

deduce the numbers of protons, neutrons and electrons present in both atoms and ions given atomic (proton)number, mass (nucleon) number and charge

BEFORE YOU START

DEVELOPING AN IDEA: NANOMACHINES

Progress in science depends not only on original thinking but also on developing the ideas of otherpeople The idea of an atom goes back over 2000 years to the Greek philosopher Demokritos About

350 years ago, Robert Boyle looked again at the idea of small particles but there was no proof JohnDalton moved a step closer to proving that atoms exist: he developed the idea that atoms of the samekind had the same weight, thinking this could explain the results of experiments on combining

different substances in terms of rearrangement of the atoms

At the beginning of the 20th century, J.J Thomson (see Figure 1.6) suggested three models of theatom His preferred model was to imagine an atom as a spherical cloud of positive charge in whichelectrons were placed A few years later, scientists working under the direction of Ernest Rutherford(see Figure 1.4) fired alpha particles (which we now know are positively charged helium nuclei) at veryhigh speeds at strips of metal only 0.0005 mm thick Most of the alpha particles went through thestrip This would fit with the idea of atoms being a cloud of charge with very little mass to deflect(change the direction of) the alpha particles But one alpha particle in every 20 000 was deflected at

an angle of more than 90° From this, Rutherford deduced that there must be something very smalland positively charged in the atom The atomic nucleus had been discovered!

In 1960 Richard Feynman (Figure 1.1) suggested that tiny machines could be made from a few

hundred atoms grouped together in clusters At the time, these ideas seemed like ‘science fiction’ Butseveral scientists took up the challenge and the science of nanotechnology was born

Without looking at the Periodic Table, make a list of the names and symbols for the elements inPeriods 1, 2 and 3 Compare your list with another learner then check to see if the symbols arecorrect

How can you deduce the formula of a simple ion (e.g a chloride ion or an aluminium ion) by

reference to the Periodic Table?

Take turns in challenging another learner to write down the formula of a simple ion Check youranswers afterwards using a textbook

Make a list of the subatomic particles in an atom giving their relative mass and relative charges aswell as their position in the atom, structure of the atom and isotopes Compare your answers withthose of another learner Were you in agreement?

Cr3+ ion If you are unsure, check your answer with someone else in the class or with a teacher.Take a photocopy of the modern Periodic Table and cross out or cut out the group numbers andperiod numbers Get another learner to select an element You then have to state in which periodand group that element belongs Take turns in doing this until you are sure that you can easilyidentify the group and period of an element

metal or metalloid (metalloids have some characteristics of both metals and non-metals) If you areboth uncertain, consult a textbook or the internet Take turns in doing this until you are sure thatyou can easily identify the position of metals, non-metals and metalloids

Ask another learner to select an element You then have to state if the element is a metal, non-Explain to another learner in terms of numbers of electrons and protons why a sodium ion has asingle positive charge but an oxide ion has a 2− charge

Explain to another learner what you know about attraction or repulsion of positive and negativecharges

In nanotechnology, scientists design and make objects that may have a thickness of only a few

thousand atoms or less Groups of atoms can be moved around on special surfaces (Figure 1.2) In thisway, scientists have started to develop tiny machines that will help deliver medical drugs to exactlywhere they are needed in the body

What else do you think nanomachines could be used for?

Every substance in our world is made up from chemical elements These chemical elements cannot bebroken down further into simpler substances by chemical means A few elements, such as nitrogen andgold, are found on their own in nature, not combined with other elements Most elements, however, arefound in combination with other elements as compounds

Every element has its own chemical symbol The symbols are often derived from Latin or Greek words.Some examples are shown in Table 1.1

The structure of an atom

Every atom has nearly all of its mass concentrated in a tiny region in the centre of the atom called thenucleus The nucleus is made up of particles called nucleons There are two types of nucleon: protons and

neutrons Atoms of different elements have different numbers of protons

Outside the nucleus, particles called electrons move around in regions of space called orbitals (seeSection 2.3) Chemists often find it convenient to use a simpler model of the atom in which electrons movearound the nucleus in electron shells Each shell is a certain distance from the nucleus at its own particular

energy level (see Section 2.3) In a neutral atom, the number of electrons is equal to the number ofprotons A simple model of a carbon atom is shown in Figure 1.3

Figure 1.3: A model of a carbon atom This model is not very accurate but it is useful for understanding what happens to

the electrons during chemical reactions

IMPORTANT

When we use a simple model of the atom we talk about shells (n = 1, n = 2, etc) and sub-shells 2s, 2p,etc In this model, the electrons are at a fixed distance from the nucleus This model is useful when wediscuss ionisation energies (Chapter 2)

When we discuss where the electrons really are in space, we use the orbital model In this model, there is aprobability of finding a particular electron within certain area of space outside the nucleus We use thismodel for discussing bonding and referring to electrons in the sub-shells

Atoms are tiny, but the nucleus of an atom is much smaller If the diameter of an atom were the size of afootball stadium, the nucleus would only be the size of a pea This means that most of the atom is emptyspace! Electrons are even smaller than protons and neutrons

a metal wire (cathode), which is heated to a high temperature when a low voltage is applied to it Atthe other end of the tube is a fluorescent screen, which glows when electrons hit it

Figure 1.5: a The beam of electrons is deflected away from a negatively charged plate and towards a positively

charged plate b The electron beam in a cathode-ray tube is deflected by an electromagnetic field The direction of

the deflection shows us that the electron is negatively charged

The electrons are given off from the heated wire and are attracted towards two metal plates, which arepositively charged As they pass through the metal plates, the electrons form a beam When theelectron beam hits the screen a spot of light is produced When an electric field is applied across thisbeam the electrons are deflected The fact that the electrons are so easily attracted to the positivelycharged anode and that they are easily deflected by an electric field shows us that:

electrons have a negative charge

electrons have a very small mass

Figure 1.6: J J Thomson calculated the charge to mass ratio of electrons He used results from experiments with

In recent years, experiments have been carried out with beams of electrons, protons and neutronsmoving at the same velocity in an electric field

Figure 1.7: A beam of protons is deflected away from a positively charged area This shows us that protons have a

positive charge

The results of these experiments show that:

a proton beam is deflected away from a positively charged plate; as like charges repel, the protonsmust have a positive charge (Figure 1.7)

an electron beam is deflected towards a positively charged plate; as opposite charges attract, theelectrons must have a negative charge

a beam of neutrons is not deflected; this shows that they are uncharged

In these experiments, huge voltages have to be used to deflect the proton beam This contrasts withthe very low voltages needed to deflect an electron beam These experiments show us that protonsare much heavier than electrons If we used the same voltage to deflect electrons and protons, thebeam of electrons would have a far greater deflection than the beam of protons This is because aproton is about 2000 times heavier than an electron

A beam of electrons is passing close to a highly negatively charged plate When the electrons passclose to the plate, they are deflected away from the plate

1

What deflection would you expect, if any, when the experiment is repeated with beams of i protons and ii neutrons? Explain your answers.

Which subatomic particle (electron, proton or neutron) would deviate the most? Explain youranswer

a

b

neutron n 1 0

Table 1.2: Comparing electrons, neutrons and protons.

Atomic (proton) number and mass (nucleon) number

The number of protons in the nucleus of an atom is called the atomic number (proton number) (Z) Every

atom of the same element has the same number of protons in its nucleus It is the atomic number thatmakes an atom what it is For example, an atom with an atomic number of 11 must be an atom of theelement sodium No other element can have 11 protons in its nucleus The Periodic Table of elements isarranged in order of the atomic numbers of the individual elements (see Appendix 1)

isotopes

Isotopes are atoms of the same element with different mass numbers

Isotopes of a particular element have the same chemical properties because they have the same number

of electrons They have slightly different physical properties, such as small differences in density or smalldifferences in mass, because they have different numbers of neutrons

We can write symbols for isotopes We write the nucleon number at the top left of the chemical symbol andthe proton number at the bottom left

Figure 1.8: The atomic structure and isotopic symbols for the three isotopes of hydrogen.

When writing generally about isotopes, chemists also name them by leaving out the proton number andplacing the mass number after the name For example, the isotopes of hydrogen can be called hydrogen-1,hydrogen-2 and hydrogen-3

Remember that in writing isotopes, mass number is the figure which is superscript

Isotopes can be radioactive or non-radioactive Specific radioisotopes (radioactive isotopes) can be used tocheck for leaks in oil or gas pipelines and to check the thickness of paper They are also used in medicine

to treat some types of cancer and to check the activity of the thyroid gland in the throat

Question

How many protons, neutrons and electrons?

In a neutral atom the number of positively charged protons in the nucleus equals the number of negativelycharged electrons outside the nucleus When an atom gains or loses electrons, ions are formed, which areelectrically charged For example:

bonded atoms Remember that positive ions are formed when one or more electrons are lost by an atom and that negative ions are formed when one or more electrons are gained by an atom.

WORKED EXAMPLE

Questions

REFLECTION

Read the paragraph in ‘Developing an idea: Nanomachines’ at the beginning of this chapter aboutRutherford’s work in discovering the nucleus Discuss these questions with another learner:

Suggest what happened to the alpha particles that went a little way from the nucleus Use ideas ofattractive or repulsive forces

Beams of protons and electrons are deflected by electric fields but neutrons are not

The atom consists of positively charged protons and neutral neutrons in the nucleus, surrounded by negativelycharged electrons arranged in energy levels (shells)

Isotopes are atoms with the same atomic number but different mass numbers They only differ in the number ofneutrons they contain

72Hf2+.Deduce the number of electrons that are present in one of these hafnium ions

[1] [1]

Nucleon number

Numbers of each subatomic particle present

Explain why a chloride ion has a single negative charge

i

ii

After studying this chapter, complete a table like this:

I can

See section

Needs more work

Almost there

Ready to move on

In this chapter you will learn how to:

understand that electrons in an atom can only exist in certain energy levels (shells) outside the nucleus

use and understand the terms shells, sub-shells, orbitals, principal quantum number (n), ground state electronicconfiguration

describe the number of orbitals making up s, p and d sub-shells and the number of electrons that can fill s, p and dsub-shells

describe the order of increasing energy of the sub-shells within the first three shells and the 4s and 4p sub-shellsdescribe the electronic configurations in each shell, sub-shell and orbital

explain the electronic configuration in terms of energy of the electrons and inter-electron repulsion

determine the electronic configuration given the atomic (proton) number and charge using either the full electronicconfiguration or shorthand electronic configuration, e.g for Fe: 1s22s22p63s23p63d64s2 (full) or [Ar]3d64s2

Ask another learner to select one of the first 20 elements in the Periodic Table Then ask them tostate the charge on the ion This time, deduce the simple electronic configuration for the ion of thiselement Take turns in doing this until you are sure that you can write the electronic configurationsfor the ions

scientists encouraged him to develop an electron microscope but it was not until 1931 that ErnstRuska and Max Knoll produced the first electron microscope

In the modern electron transmission microscope, a beam of electrons is accelerated in an electric fieldand focused on a specimen by an electromagnetic lens Some electrons are transmitted through thespecimen, others are absorbed and reflected as various types of radiation The transmitted andreflected radiation can be absorbed onto a fluorescent screen to give an image (Figure 2.1)

The electron microscope can be used to calculate the size of pores in polymers and investigate thesurfaces of metals (see Figure 2.2) and other materials The radiation reflected in a scanning electronmicroscope can be used to identify chemical elements

Why must the specimens studied using an electron microscope have dry surfaces?

Why might it be useful to be able to look at a substance in more detail?

In Section 1.2 we saw that electrons are arranged outside the nucleus in energy levels or quantum shells.These principal energy levels or principal quantum shells (symbol n) are numbered according to how farthey are from the nucleus The lowest energy level, n = 1, is closest to the nucleus, the energy level n = 2

is further out, and so on The electrons in quantum shells further away from the nucleus have more energyand are held less tightly to the nucleus

The arrangement of electrons in an atom is called its electronic structure or electronic configuration.The electronic configurations of lithium, neon and chlorine are shown in Figure 2.3, together with a

shorthand way of writing this structure

Figure 2.3: The simple electronic structures of lithium, neon and chlorine The nuclei of the atoms are not shown.

Table 2.1 shows the number of electrons in each of the principal quantum shells (principal energy levels)for the first 11 elements in the Periodic Table

1

sulfur; the atomic number of sulfur, Z = 16

a

magnesium, Z = 12 fluorine, Z = 9 potassium, Z = 19 carbon, Z = 6

b

c

d

e

Ionisation energy, IE

By firing high-speed electrons at atoms, scientists can work out how much energy has to be supplied toform an ion by knocking out one electron from each atom

Ionisation energies are measured under standard conditions The general symbol for ionisation energy is IE.Its units are kJ mol−1

The symbol for the first ionisation energy is IE1 Using calcium as an example:

1st ionisation energy: Ca(g) → Ca+(g) + e−

IE1 = 590 kJ mol−1

If a second electron is removed from each ion in a mole of gaseous 1+ ions, we call it the second ionisation energy, IE2 Again, using calcium as an example:

IMPORTANT

When you write equations for successive ionisation energies, remember that:

the atoms and ions are in the gaseous state

the charge on the ion on the right-hand side gives the number of the ionisation energy (1st, 2nd,3rd, etc.)

The successive ionisation energies for the first 11 elements in the Periodic Table are shown in Table 2.2.The data in Table 2.2 shows us that:

For each element, the successive ionisation energies increase This is because the net positive charge

on the ion gets greater as each electron is removed As each electron is removed there is a greaterattractive force between the positively charged protons in the nucleus and the remaining negativelycharged electrons Therefore, more energy is needed to overcome these attractive forces

There is a big difference between some successive ionisation energies For nitrogen, this occurs

between the 5th and 6th ionisation energies For sodium, the first big difference occurs between the1st and 2nd ionisation energies These large changes indicate that for the second of these two

ionisation energies the electron is being removed from a principal quantum shell closer to the nucleus.For example, for the 5th ionisation energy of nitrogen, the electron being removed is from the 2nd principalquantum shell For the 6th ionisation energy of nitrogen, the electron being removed is from the 1st

Factors that influence ionisation energy

Four factors that influence ionisation energy are:

The size of the nuclear charge: as the atomic number (number of protons) increases, the positivenuclear charge in the nucleus increases The bigger the positive charge, the greater the attractiveforce between the nucleus and the electrons So, more energy is needed to overcome these attractiveforces if an electron is to be removed

In general, ionisation energy increases as the proton number increases

Distance of outer electrons from the nucleus: the force of attraction between positive and negativecharges decreases rapidly as the distance between them increases So, electrons in shells further awayfrom the nucleus are less attracted to the nucleus than those closer to the nucleus

In general, the further the outer electron shell is from the nucleus, the lower the ionisation energy.Shielding effect of inner electrons: as all electrons are negatively charged, they repel each other.Electrons in full inner shells repel electrons in outer shells Full inner shells of electrons prevent theouter electrons feeling the full nuclear charge This is called shielding The greater the shielding ofouter electrons by the inner electron shells, the lower the attractive forces between the nucleus andthe outer electrons See Section 2.2 for information on sub-shells and how they affect shielding

In general, the ionisation energy is lower as the number of full electron shells between the outerelectrons and the nucleus increases

Spin-pair repulsion: electrons in the same atomic orbital in a sub-shell (see Section 2.3) repel eachother more than electrons in different atomic orbitals This increased repulsion makes it easier toremove an electron So first ionisation energy is decreased

Interpreting successive ionisation energies

Figure 2.5 shows a graph of successive ionisation energies against the number of electrons removed forsodium A logarithmic scale (to the base 10) is used because the values of successive ionisation energieshave such a large range

We can deduce the following about sodium from Figure 2.5:

The first electron removed has a low first ionisation energy, when compared with the rest of the data

It is very easily removed from the atom It is therefore likely to be a long way from the nucleus andwell shielded by inner electron shells

The second electron is much more difficult to remove than the first electron There is a big jump in thevalue of the ionisation energy This suggests that the second electron is in a shell closer to the nucleusthan the first electron Taken together, the 1st and 2nd ionisation energies suggest that sodium hasone electron in its outer shell

The 10th and 11th electrons have extremely high ionisation energies, when compared with the rest ofthe data This suggests that they are very close to the nucleus There must be a very great force ofattraction between the nucleus and these electrons and there are no inner electrons to shield them.The large increase in ionisation energy between the 9th and 10th electrons confirms that the 10thelectron is in a shell closer to the nucleus than the 9th electron

The second principal quantum level, n = 2, can hold a maximum of 8 electrons: 2 electrons in the ssub-shell and 6 electrons in the p sub-shell

shell, 6 electrons in the p sub-shell and 10 electrons in the d sub-shell

The third principal quantum level, n = 3, can hold a maximum of 18 electrons: 2 electrons in the s sub-You will also notice from Figure 2.7 that the order of the sub-shells in terms of increasing energy does notfollow a regular pattern of s then p then d after the element argon The order of sub-shells after argonappears to overlap The next element after argon is potassium Potassium’s outer electron is in the 4s, not

in the 3d, sub-shell The first element with an electron in the 3d sub-shell is element 21, scandium

Figure 2.7: The sub-shells for the first four principal quantum shells The maximum number of electrons that are allowed

in each sub-shell is: s = 2 electrons, p = 6 electrons, d = 10 electrons

When high-speed electrons hit gas particles at low pressure, coloured lines are seen through an instrumentcalled a spectroscope (see Figure 2.4) The letters s, p and d come from the terms used to describe theselines: ‘s’ for ‘sharp’, ‘p’ for ‘principal’ and ‘d’ for ‘diffuse’

Atomic orbitals

Each sub-shell contains one or more atomic orbitals

An atomic orbital is a region of space around the nucleus of an atom that can be occupied by one or twoelectrons

As each orbital can only hold a maximum of two electrons, the number of orbitals in each sub-shell mustbe:

as the 1s orbital in the first quantum shell They are both spherical, but electrons in the 2s orbital havemore energy than electrons in the 1s orbital There are three 2p orbitals in the second quantum shell Each

of these has the same shape The shape is like an hourglass with two ‘lobes’ The three sets of ‘lobes’ are

arranged at right angles to each other along the x, y and z axes Hence the three 2p orbitals are named

2px, 2py and 2pz The three 2p orbitals have the same energy as each other There are also three 3p

orbitals in the third quantum shell Their shapes are similar to the shapes of the 2p orbitals but, again, theirenergies are higher

Figure 2.8: Representations of orbitals (the position of the nucleus is shown by the black dot) a s orbitals are spherical b

p orbitals, px, py and pz , have ‘lobes’ along the x, y and z axes.

Figure 2.9: The shape of a dz2 orbital

The d orbitals are more complex in shape and arrangement in space In 1925 Louis de Broglie suggestedthat electrons behaved like waves This led to the idea of electron probability clouds The electron

probability cloud for one type of d orbital has two pieces: it is like a modified p orbital with a ring aroundthe middle (Figure 2.9) You will not need to know the d-orbital shapes at AS Level, but you will for A Levelwhen studying the transition elements (see Chapter 24)

Filling the shells and orbitals

The most stable electronic configuration of an atom is the one that has the lowest amount of energy Theorder in which the sub-shells are filled depends on their relative energy The sub-shell with the lowestenergy, the 1s, is therefore filled first, followed by those that are successively higher in energy As we

noted in Figure 2.7, the order of the sub-shells in terms of increasing energy does not follow a regular