chapter 03 the periodic 0 table

chapter 03

chapter outcomes

key knowledge

00

After completing this chapter, you should be able to:

The periodic table

03

chhaptter outtcomesAfter completing this chapter, yoube able to:

• historical development of the periodic table from Mendeleev to Seaborg

• trends and patterns of properties within the periodic table: atomic number, types of compounds formed, metallic/non-metallic character, chemical reactivity of elements

• relate the development of the periodic table to the experimental work and hypotheses of chemists from Mendeleev to Seaborg

• explain the periodic variation of properties of the elements in terms of electronic confi gurations

• map the s-, p-, d- and f-blocks on a blank table• given its electronic confi guration, place an element

in the periodic table• locate metals and non-metals in the periodic table• give an indication of the trend in the chemical

reactivities of the elements • explain why elements and compounds are both

classifi ed as pure substances• explain the trends and patterns in the following

atomic properties of the main group elements: atomic number and radius, ionisation energy, electronegativity and types of compounds formed

• use electronic confi guration and atomic properties to predict the reactivity of the main group elements.

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03333333303003The periodic table

37

3.1

Why is the periodic table important?Chapter 1 dealt with the beginnings of chemistry, from the random appli-cations of chemical technology to the optimistic experimental aims of the alchemists and on to the more systematic search for new elements by the early chemists. The alchemists and chemists gathered a large amount of chemical information and we saw how chemists such as Boyle, Lavoisier, Dalton and others tried to make sense of what they were learning. Against this background, Mendeleev’s periodic table was an inspired development. Not only did it organise the information gathered on the basis of patterns of chemical properties and behaviour, but it also proved to be a useful tool for predicting and guiding future developments in chemistry. However, it did not adequately explain why substances had certain properties and behaved as they did.

Chapter 2 told some of the stories of the scientists who carried out research to reveal the nature of atoms: the existence of subatomic particles, their charge and mass, the way these particles are arranged in an atom and the way they behave. This search also illustrated the way in which scientists such as Thomson, Rutherford, Bohr and Schrödinger developed successive models of the atom. Each model was consistent with some or all of existing

Figure 3.1Our knowledge of the properties of the elements and the reactions they undergo helps chemists to select appropriate chemicals for countless applications.

!‘The essence of science—ask an impertinent question and you are on the way to fi nding the pertinent answer.’

Dr Jacob Bronowski

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38 030000000333030003The periodic table

38hTh

!‘Halogen’ comes from the Greek halos ‘salt’ and ‘gen’ create. A compound formed between a halogen and a metal is called a ‘salt’. Common salt used in cooking is sodium chloride.

knowledge but subject to review when new experimental data highlighted its limitations. Science is a collaborative activity and scientists must be ready to subject their experiments and ideas to the scrutiny of their peers.

The ideas from Chapter 1 combined with the understanding that comes from Chapter 2 has made the periodic table the most useful reference available to chemists. It minimises the need to memorise isolated facts and provides a framework with which to organise our knowledge. Knowing the properties of particular elements and trends within the table, chemists can comprehend what would otherwise be an overwhelming collection of disorganised experimental observations. With the table at hand, chemists can extrapolate from known facts to the unknown.

3.2

Why are element properties periodic?Mendeleev discovered sets of elements with similar chemical properties. As he looked at the elements in order of increasing atomic mass, he saw that they changed from metals to non-metals, back to metals again, and so on, according to a regular pattern. Similar patterns existed for other properties. Mendeleev proposed his periodic law: The properties of elements vary periodically with their atomic weights.

The number of protons in an atom, or atomic number, is now known to be the fundamental difference between atoms of different elements, rather than an element’s atomic weight. As a result, the modern table has elements arranged in order of increasing atomic number.

To understand the reason for the periodicity of element properties, let us look at two groups of elements that Mendeleev recognised as being similar.

The fi rst group is the alkali metals. The elements in this group (lithium, sodium, potassium, rubidium and caesium) are all relatively soft metals and are highly reactive with water and oxygen.

Consider their electronic confi gurations:

Li 1s22s1

Na 1s22s22p63s1

K 1s22s22p63s23p64s1

Rb 1s22s22p63s23p63d104s24p65s1

Cs 1s22s22p63s23p63d104s24p64d105s25p66s1

These elements each have similar outer-shell electronic confi gurations—all have one electron in an s-subshell.

The second group we will look at is the halogens. Fluorine, chlorine, bromine and iodine are all coloured and are also highly reactive, especially fl uorine and chlorine. Their electronic confi gurations are:

F 1s22s22p5

Cl 1s22s22p63s23p5

Br 1s22s22p63s23p63d104s24p5

I 1s22s22p63s23p63d104s24p64d105s25p5

These elements also have similar outer-shell electronic confi gurations; in this case each element has s2p5 as the outer shell. The electrons in the outer shell of an atom are known as the valence electrons.

review

Periodicity refers to the repeated pattern of similar chemical properties as elements are arranged in order of increasing atomic number. Elements with similar chemical properties are placed in the same group of the periodic table.

ATOMIC STRUCTURE

• Periodic table and electron confi guration

E

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39

Li3

Na11

K19

Rb37

Cs55

Fr87

Be4

H1

He2

Mg12

Ca20

Sr38

Ba56

B5

Al13

Ga31

In49

Tl81

C6

Si14

Ge32

Sn50

Pb82

N7

P15

As33

Sb51

Bi83

O8

S16

Se34

Te52

Po84

F9

Cl17

Br35

I53

At85

Ne10

Ar18

Kr36

Xe54

Rn86

Ra88

Sc21

Ti22

Zr40

Hf72

Ce58

Th90

Pr59

Pa91

Nd60

U92

Pm61

Np93

Sm62

Pu94

Eu63

Am95

Gd64

Cm96

Tb65

Bk97

Dy66

Cf98

Ho67

Es99

Er68

Fm100

Tm69

Md101

Yb70

No102

Lu71

Lr103

V23

Nb41

Ta73

Cr24

Mo42

W74

Mn25

Tc43

Re75

Fe26

Ru44

Os76

Co27

Rh45

Ir77

Ni28

Pd46

Pt78

Cu29

Ag47

Au79

Zn30

Cd48

Hg80

I II III IV V VI VII VIII

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

s1 s2

Period

2

3

4

5

6

7

1

s2p1 s2p2 s2p3 s2p4 s2p5 s2p6

Transition elements

Lanthanides

Actinides

Y39

La57

Ac89

Rf104

Db105

Sg106

Bh107

Hs108

Mt109

Ds110

Rg111

Uub112

Uut113

Uuq114

Uup115

Uuh116

Group

New number:

Old number:

Outer-shellelectronic

configuration:

!The recurrence of the same outer-shell electronic confi guration is responsible for the elements in a group having similar chemical properties.

The arrangement of electrons in atoms is responsible for the periodicity of element properties.

Melting temperature, electrical conductivity, formulas of compounds formed when two or more elements react, and many other properties depend on the way electrons are arranged in the atoms involved. In particular, the outer-shell electronic confi guration of an atom is especially important; these are the electrons that are generally involved in forming chemical bonds. Similar electronic confi gurations occur regularly as the atomic number of atoms increases; this causes periodicity in chemical properties of elements. Mendeleev’s periodic law can be restated in modern terms:

Variations of the chemical properties of elements across a period and similarities down a group are all associated with the electronic confi gurations of their atoms.

Patterns in electronic structureWe now regard the periodic table as an arrangement of elements showing similarities in electronic confi gurations. Many versions of the periodic table have been designed to emphasise the relationships that exist between elements. No one version is ideal, each having disadvantages and advantages over others.

!Two other versions of the periodic table are shown on page 52.

Figure 3.2This form of the periodic table is in common use. Colour is used to distinguish the blocks of elements.

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The ‘long form’ of the table shown in Figure 3.2 is in common use today. This table has the following features: Vertical columns are called groups. Groups contain elements with similar

outer-shell electronic confi gurations. For instance, atoms of the alkali metals in the fi rst group each have a half-fi lled s-orbital (s1) in their outer shell.

Groups are numbered 1–18. Alkali metals, with s1 outer shells form group 1; the halogens, with s2p5 outer shells, form group 17.

Horizontal rows are called periods. Periods are numbered 1–7 and each contains elements with electrons in the same outer shell. The number of the period is the same as the number of the outer shell.

For example, the outer shell of elements such as magnesium and chlorine is the third shell and both these elements are in period 3:

Mg 1s22s22p63s2

Cl 1s22s22p63s23p5

Similarly, the elements in period 5 all have their outer-shell electrons in the fi fth shell.

Blocks of elementsThe periodic table can be regarded as having four main blocks of elements with the elements in each block all fi lling the same type of subshell: The s-block contains the elements in group 1 (alkali metals) and group

2 (alkaline earth metals). These elements have a half-fi lled or fi lled s-subshell, that is s1 or s2, as the highest energy subshell.

The p-block contains elements in groups 13–18. A p-subshell is the highest-energy subshell of elements in this section of the table. The elements have outer-shell confi gurations from s2p1 to s2p6.

Elements in the s- and p-blocks are often referred to as main group elements.

Although helium has an electronic confi guration of 1s2, it is usually located in group 18 because its atoms are unreactive—like the other atoms in group 18, which have an s2p6 outer shell. Deciding where in the periodic table to place hydrogen is also diffi cult. Although hydrogen atoms have an s1 outer shell, hydrogen behaves quite differently from group 1 metals. To resolve these complications, we have placed hydrogen on its own rather than in group 1 but you will notice that some periodic tables place it in group 1. The d-block, which contains the elements known as transition metals,

is located between the s- and p-blocks. It contains elements in which the fi ve orbitals, with a maximum of 10 electrons, of a d-subshell are progressively being fi lled. We have seen that d-subshells are fi lled only after the s-subshell of the next shell has been fi lled. As a consequence, the transition metals have outer-shell confi gurations from d1s2 to d10s2. Transition elements include many metals of importance to modern society (Figure 3.3).

The f-block contains the lanthanides and actinides. The lanthanides are a set of 14 elements with atomic numbers 58–71. The actinides have atomic numbers 90–103. The seven orbitals, able to contain a maximum of 14 electrons, in the 4f-subshell are progressively being fi lled in the lanthanides and the 5f-orbitals are being fi lled in the actinides.

Although not classifi ed as a separate block, it is worth commenting on the group of elements with atomic numbers from 93 on. Uranium is the last element to be been found naturally on Earth. The transuranic (higher atomic number than uranium) elements have been created in laboratories, the fi rst identifi ed as element number 93 and called neptunium in 1940. Credit for a major role in producing many of these elements is given to Glen Seaborg and his colleagues.

Figure 3.3 This stainless steel sculpture by Bert Flugelman in Martin Place, Sydney, is made from the transition metals chromium and iron.

!All the elements in a period have the same number of occupied electron shells.

chemfactA metallic form of hydrogen is believed to exist in the high-pressure core of the planet Jupiter.

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Glenn Seaborg, with a team of scientists at the University of California in Berkeley, used a particle accelerator called a cyclotron to bombard uranium with hydrogen nuclei. In February 1941, they experimentally confi rmed the existence of element 94 among the products of this bombardment. The element was named plutonium, symbol Pu.

Research showed that plutonium-239 could readily undergo nuclear fi ssion—the splitting of the nucleus into two or more smaller fragments. Like nuclear fusion (the combination of two nuclei to form a larger nucleus), fi ssion reactions are accompanied by the release of large amounts of energy. The United States Government recognised the weapons potential of the discovery and assigned Seaborg to the Manhattan Project. It took 4 years of research to make the plutonium bomb that was detonated over Nagasaki on 9 August 1945. Plutonium is now an integral part of many modern nuclear weapons. It is also one of the most toxic substances known.

On his return to Berkeley in 1944, Seaborg continued his experiments with the transuranic elements. Elements 95 and 96 were produced by bombarding plutonium and were named americium and curium (after Marie Curie). These new elements

were then also bombarded in a cyclotron. Elements 97 and 98 were discovered in 1949 and 1950 and named berkelium and californium respectively.

The discoveries of elements 99, 100, 101 and 103 were all credited to Seaborg’s team. Then in 1964, Soviet scientists led by Georgii Flerov claimed to have produced element 104 and, 3 years later, element 105.

The Californian group produced the elements by their own methods and gave them different names. A dispute arose and, to resolve it, element 104 and those higher became known by systematic names.

The International Union of Pure and Applied Chemistry (IUPAC) formally registers the names of chemical elements and compounds. IUPAC requires experimental demonstration, beyond reasonable doubt, of a nuclide with an atomic number Z not identifi ed before, existing for at least 10–14 s.

For a time, element 104 was called unnilquadium, symbol Unq, but is now known as rutherfordium. Similarly, elements 106, seaborgium, and 107, bohrium, are named after key contributors to our understanding of the atom. Elements 108–112 have been identifi ed by scientists working in Darmstadt, Germany, under the leadership of Peter Armbruster. The name of element 110, discovered in 1994, was confi rmed by IUPAC in August 2003 as darmstadtium (Ds). On 1 November 2004, element 111 was offi cially named roentgenium (Rg), to honour Wilhelm Conrad Roentgen who discovered X-rays in 1895. So widespread were the benefi ts of X-rays that, in 1901, Roentgen was awarded the fi rst Nobel Prize in Physics.

Many of these elements are prepared in minute quantities and may be short-lived, but their formation enables scientists to test and refi ne their understanding of the structure of the nucleus.

The discovery of the transuranic elements

extension

E1 To which block of the periodic table does: a plutonium belong? b roentgenium belong?

E2 a Give the name and chemical symbol for each of the elements 99, 100, 101 and 102. b In honour of whom were each of the elements in part a named?

questions ?

Gl S b ith t f i ti t t th U i it f C lif i

Figure 3.4 Glenn Seaborg and colleague Darleane Hoffman. Research teams led by Seaborg discovered many of the transuranic elements. Darleane Hoffman discovered the existence of plutonium-244 and studied the aqueous chemistry of element 105. Glenn Seaborg was awarded the Nobel Prize in Chemistry in 1951 for his contributions to the understanding of transuranic elements.

Figure 3.5 A replica of Little Boy, the uranium-235 bomb that was dropped on Hiroshima on 6 August 1945.

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3.2

The modern periodic table lists elements in order of increasing atomic number.

The common long form of the periodic table contains vertical columns or groups numbered 1–18. Elements in the same group have the same number of valence electrons.

Horizontal rows are called periods. Elements in the same period have valence electrons in the same outer shell.

The periodic table is also divided into four blocks labelled s, p, d and f. These refl ect the subshell occupied by the valence electrons of the elements in the block.

Common names are given to elements in some of the groups in the periodic table. Elements in:• group 1 elements are the alkali metals• group 17 elements are the halogens• group 18 elements are the noble gases• groups 3–12 (the d-block elements) are the transition elements.

summary

1 Use the periodic table (Figure 3.2) to answer the following questions.

a In which group of the periodic table will you fi nd the following? i B iv Ar ii Cl v Si iii Na vi Pb b In which period of the periodic table will you fi nd the following? i K iv H ii F v U iii He vi P c What is the name, symbol and electronic confi guration of: i the second element in group 14? ii the second element in period 2? iii the element which is in group 18 and period 3? 2 Use the periodic table to predict the number of valence

electrons for each of the following: a Ba d Te b Br e Ge c Cs 3 Suppose you were familiar with the properties, reactions

and compounds formed by the following elements: sodium,

chlorine, magnesium, argon and sulfur. Which element would you use to predict the properties of the following?

a krypton d potassium b selenium e barium c fl uorine 4 Name an element with similar properties to those of: a potassium c selenium b barium d boron 5 In which block in the periodic table are elements with the

following ground state electronic confi gurations located? Name the period and group, where appropriate:

a 1s22s22p63s2 c 1s22s22p63s23p63d54s2

b 1s22s22p2 d 1s22s22p63s23p4

6 Name the period and group of each of the following elements: a the element with a +3 charged ion that has an electronic

confi guration of 1s22s22p6

b the element with an excited atom that has an electronic confi guration of 1s22s22p63s13p6

c the element with a −2 charged ion that has an electronic confi guration of 1s22s22p63s23p6

7 Why are elements arranged in the periodic table in order of atomic number rather than relative atomic weight?

key questions

3.3

Trends in propertiesThe periodic variation in the properties of elements refl ects the periodic variation in their electronic confi gurations. This variation is most obvious if we look at how the properties change as we move from left to right across a period. Such trends refl ect the effect of successively adding an electron as we move from one element to the next.

There are similarities between elements in a group but signifi cant differences also occur.

Although all elements within a group have the same number of electrons in their outer shells, the atoms become larger as we move down a group and this also has an effect on the properties of the elements.

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We can look at the properties of elements in two ways.1 The properties that relate to individual atoms of the element. Electronic

confi guration is one such atomic property.2 The properties that refl ect the way atoms or groups of atoms interact

with each other. Such interactions would include those between atoms in a piece of aluminium, those between oxygen atoms in a sample of oxygen gas and those between aluminium and oxygen when a piece of aluminium is left in the air. These three examples involve three very different types of interactions. They will be covered in greater detail in Chapters 5–7.

We will look at the atomic properties of elements fi rst as these determine the way atoms interact and react.

Atomic propertiesThe properties of radius, ionisation energy and electronegativity all depend on the strength of the attraction between the outer-shell electrons and the nucleus. In general terms, this attraction will depend on: the positive charge that attracts the outer-shell electrons the distance of the electrons from the nucleus.

Figures 3.6–3.8 summarise the relevant data for trends in these three atomic properties.

Atomic radius decreases

Ato

mic

rad

ius

incr

ease

s

1

Li

Na

K

Rb

Cs

Be

Mg

Ca

Sr

Ba

152

186

227

247

265

113

160

147

215

217

B

Al

Ga

In

Tl

88

143

122

163

170

C

Si

Ge

Sn

Pb

77

H

37

117

122

140

173

N

P

As

Sb

Bi

70

110

121

141

155

O

S

Se

Te

Po

66

104

117

143

167

F

Cl

Br

I

At

64

99

114

133

143

2 13 14 15 16 17

Figure 3.7 Variation in fi rst ionisation energy with atomic number

Li1.0Na0.9K

0.8Rb0.8Cs0.8Fr0.7

Be1.6Mg1.3Ca1.0Sr1.0Ba0.9Ra0.9

H2.1

B2.0Al1.6Ga1.8In

1.8Tl

2.0

C2.6Si

1.9Ge2.0Sn2.0Pb2.3

N3.0P

2.2As2.2Sb2.1Bi2.0

O3.4S

2.6Se2.6Te2.1Po2.0

F4.0Cl3.2Br3.0

I2.7At2.2

Electronegativity increases across a period.

Electronegativityincreases upa group.

Figure 3.8 Trends in electronegativity in the periodic table.

Figure 3.6 Relative size of atoms of selected main group elements. The atomic radius is given in picometres. A picometre is 10−12 m.

chemfactElectronegativity is a measure of the ability of an atom to attract an electron towards itself. Non-metals tend to have a higher electronegativity than metals.

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11+ 17+

Core charge feltby outer-shellelectron = +1


sodium chlorine

3+ 9+



lithium fluorine

Figure 3.9Electron shell diagrams of lithium, sodium, fl uorine and chlorine. For simplicity, the electrons have been represented in shells rather than subshells.

The electrons in atoms of elements in the same period are located in the same outer shell. As you move across the period, the number of protons in the nucleus increases, causing the number of electrons in the atom to increase. Consider the atoms of two of the period 2 elements, lithium and fl uorine, and the period 3 elements, sodium and chlorine, shown in Figure 3.9: The outer-shell electron of a lithium atom is attracted towards the three

protons in the atom’s nucleus. However, the outer-shell electron does not feel the full attraction of the nuclear charge as the two inner-shell electrons tend to act as a shield. We can think of the outer electron as experiencing a core charge of +1.

The outer-shell electrons in a fl uorine atom experience a core charge of +7 as the two inner-shell electrons shield the outer-shell electrons from the nine protons.

Similarly, the outer-shell electron of a sodium atom is attracted towards the 11 protons in the atom’s nucleus. However, the outer-shell electron does not feel the full attraction of the nuclear charge as the 10 inner-shell electrons tend to act as a shield. We can think of the outer electron as experiencing a core charge of +1.

As with fl uorine, the outer-shell electrons in a chlorine atom experience a core charge of +7.We can see that:

Moving down a group, the core charge remains constant but the number of electron shells increases. The outer-shell electrons are therefore held less strongly.

The electrons in atoms of elements in the same period are located in the same outer shell. Moving across a period, as the core charge increases, the outer-shell electrons are pulled more and more strongly towards the atom’s nucleus.

These changes explain most of the trends in the physical properties of the elements, as shown in Table 3.1.

!Atoms do not have sharply defi ned boundaries and so it is not possible to directly measure their radii. One method of obtaining atomic radii, and therefore an indication of atomic size, is to measure the distance between nuclei of atoms in molecules. For example, in a hydrogen molecule (H

2) the two nuclei are 74 pm apart.

The radius of each hydrogen atom is assumed to be half of that distance, i.e. 37 pm.

!Core charge can also be calculated as ‘atomic number’ minus ‘all electrons except those in the outer shell’, e.g. for

17Cl, 1s22s22p63s23p5, core

charge = 17 − 10 = +7.

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TABLE 3.1 Trends within the periodic tableTrend Explanation

Atomic radius increases down a group.

Electrons occupy most of the volume of an atom. Potassium, electronic confi guration of 1s22s22p63s23p64s1, is much larger than lithium, electronic confi guration 1s22s1.

Atomic radius decreases across a period.

The size of atoms decreases. The increasing positive charge of the nucleus pulls the outer-shell electrons closer, causing the volume of an atom to reduce (Figure 3.9).

First ionisation energy decreases down a group.

As the atoms become larger, their outer electrons are further from the nucleus. The energy required to extract the outermost electron from an atom (the fi rst ionisation energy) decreases.

First ionisation energy increases across a period.

As the strength of attraction between the outer electrons and the nucleus increases, the energy required to remove the outermost electron from an atom increases.

Electronegativity decreases down a group.

As the outer electrons become more distant, electrons are more weakly attracted to an atom (Figure 3.8).

Electronegativity increases across a period.

The electron-attracting ability of atoms increases as the pull on the outer electrons increases (Figure 3.7).

Metallic and non-metallic characterElements on the right side of the table are non-metals (Figure 3.2). Other elements are metals. As we move from left to right across a period, the elements become less metallic in character and exhibit more of the properties typical of non-metals.

There is also a variation in metallic character within groups. This is perhaps most obvious in group 14. Carbon, the fi rst element in the group, is a non-metal, whereas tin and lead are metals.

Some elements, such as germanium, silicon, arsenic and tellurium which are located partway across periods, display both metallic and non-metallic properties. Such elements are called metalloids.

Chemical reactivity of the elementsReactivity of metals

The way metals react with water can give us an indication of their relative reactivity. Table 3.2 describes the reaction of some of the groups 1 and 2 metals with water.

review

Metals are usually solid at room temperature and tend to be shiny, conduct electricity and heat and can be hammered into shape. Non-metals are usually gases, liquids or low melting temperature solids at room temperature. They are dull in colour, poor conductors of electricity and heat and are brittle.

!Silicon is the second most abundant element in the Earth’s crust after oxygen. Compounds of silicon and oxygen make up most of the sand, rocks and clay on Earth. The element itself is a semiconductor, i.e. it conducts electricity to a small extent. We make use of the semiconducting properties of silicon in electronics.

Period Group Element Reaction with water

3 1 Sodium Reacts vigorously, producing enough energy to melt the sodium which fi zzes and skates on the water surface

4 1 Potassium Reacts violently, making crackling sounds as the heat evolved ignites the hydrogen produced by the reaction

5 1 Rubidium Explodes violently on contact with water

3 2 Magnesium Will not react with water at room temperature but will react with steam

4 2 Calcium Reacts slowly with water at room temperature

TABLE 3.2 Reaction of selected groups 1 and 2 metals with water. In each case, a reaction results in the formation of hydrogen gas

!Group 1 metals must be handled with great care. They are so reactive that they need to be stored under oil to prevent contact with moisture in the atmosphere.

Those observations indicate that, for the fi ve metals in Table 3.2: those in group 1 are more reactive in water than those in group 2 down a group, the reactivity of the metal in water increases.

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These generalisations are consistent with the results of experimental investigations of the reactions of other metals.

In general, the reactivity of metals increases going down a group and decreases across a period. This trend in reactivity can be explained in terms of the changes in atomic size and ionisation energy described in Table 3.1. When metals react, their atoms tend to form positive ions by donating one or more of their valence electrons to other atoms. The metal atoms with lowest fi rst ionisation energy will require least energy to donate electrons and so tend to be most reactive. Therefore, the most reactive metals are found in the bottom left-hand corner of the periodic table. (The reaction between metals and non-metals is discussed in Chapter 6.)

Reactivity of non-metals

Consider the halogens, the elements of group 17 of the periodic table. Like the elements in group 1, these elements are very reactive. We can get an indication of the relative reactivity of chlorine, bromine and iodine by reacting each of them with aqueous solutions of potassium bromide (KBr) and potassium iodide (KI).

The results of one such experiment is shown in Table 3.3.

Figure 3.10 a Potassium stored under oil. b The reaction of potassium with water releases considerable energy. This excites the outer-shell electrons of potassium, which results in the lilac fl ame colour characteristics of potassium.

These results indicate that, in this experiment, chlorine is the most reactive and iodine the least reactive.

This is consistent with other experimental data that shows that the reactivity of the halogens decreases from fl uorine in period 2 to iodine in period 5.

This trend in reactivity can be explained in terms of the changes in atomic size and electronegativity described in Table 3.1. When non-metals like the halogens react, their atoms tend to form negative ions by accepting electrons from other atoms. Fluorine, being the most electronegative of the group, has the highest attraction for electrons and so forms negative ions most readily.

The noble gases The elements of group 18 are known as the noble gases. Helium, neon, argon, krypton and xenon are all very unreactive gases. They are known as the noble gases in recognition of their unreactive properties. The heaviest member of this family of gases is radon, a radioactive gas fi rst discovered in 1900.

The elements have very low melting and boiling temperatures and are all gases at normal temperatures. Helium boils at −269°C, just 4° above the lowest temperature possible.

The lack of reactivity of these elements arises from the arrangement of electrons in their atoms (Table 3.4). With the exception of helium, the noble gases have outer-shell electronic confi gurations of s2p6, an arrangement

Period Element Reaction with KBr solution Reaction with Kl solution

3 Chlorine Reacts to form a brown solution

Reacts to form a brown solution

4 Bromine No reaction Reacts to form a brown solution

5 Iodine No reaction No reaction

TABLE 3.3

!Chlorine and bromine are both very toxic and corrosive. They must be handled with extreme care using suitable protective clothing, gloves and eye and face protection. Iodine is also harmful by inhalation and in contact with skin and precautions are needed when handling it.

chemfactWilliam Ramsay was awarded the Nobel Prize in Chemistry in 1904 for his discovery of the noble gases.

a

b

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Figure 3.11 Argon gas is used during welding to provide an unreactive atmosphere around the hot metal.

which appears to be particularly stable. Helium has two electrons only. The fi rst electron shell has a maximum capacity of two electrons so, in helium, this fi rst shell is full. This fi lled shell confers stability to the helium atom in a similar manner to that of the other noble gases.

For some years, the unreactive nature of the noble gases led many chemists to believe that these elements would form no compounds at all. However, in 1962 a British–Canadian chemist, Neil Bartlett, mixed xenon with the highly reactive compound platinum hexafl uoride (PtF6) and produced a yellow solid, xenon hexafl uoroplatinate (XePtF6). Since then, fl uorine compounds of xenon, krypton and radon have been prepared, including xenon difl uoride (XeF2) (Figure 3.12), krypton tetrafl uoride (KrF4) and radon tetrafl uoride (RnF4).

Element Electronic confi guration

Helium 1s2

Neon 1s22s22p6

Argon 1s22s22p63s23p6

Krypton 1s22s22p63s23p63d104s24p6

Xenon 1s22s22p63s23p63d104s24p64d105s25p6

Radon 1s22s22p63s23p63d104s24p64d104f145s25p65d106s26p6

TABLE 3.4 Electronic confi guration for the noble gases

Figure 3.12 Xenon difl uoride crystals. For most of the 20th century, many chemists thought this compound could not exist.

Symbol Abundance (% in theatmosphere

Uses

He 0.000 5 Mixed with argon as an inert gas blanket for arc weldingInfl ates balloons and airshipsMixed with oxygen for a breathing mixture for deep-sea diving

Ne 0.002 A fi lling in neon signs

Ar 0.93 An inert fi lling in light bulbs, sometimes with kryptonMixed with helium as an inert gas blanket for arc welding

Kr 0.000 1 Mixed with argon in lamps

Xe 0.000 08

Rn Trace Radioactive gas

TABLE 3.5 The abundance and uses of the noble gases

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3.4

CompoundsSo far we have considered substances consisting of only one kind of element. These are all pure substances. All other pure substances are called compounds. Compounds are formed when atoms of two or more elements chemically combine in fi xed proportions. Each compound has its own characteristic set of properties, which are quite different from those of the elements it contains.

You would know the compound sodium chloride as common salt. Common salt is found worldwide and always contains 39.3% sodium and 60.7% chlorine, by mass (Figure 3.13).

Table 3.6 lists some of the properties of sodium, chlorine and sodium chloride. Each of these properties tells us something about the particles present and the strength of the forces between the particles. This will be

3.3

Across a period in the periodic table, or down a group, there is a periodic variation in the properties of the elements.

Moving down a group, the number of occupied electron shells of the elements’ atoms increases and therefore so does the radius of the atoms. First ionisation energy and electronegativity decrease down a group and the elements become more metallic in character.

Across a period from left to right, there is an increase in core charge. The strength of attraction of the nucleus for electrons increases, pulling the outer-shell electrons closer and so

reducing the radius of the atom. Because of the increasing nuclear pull, fi rst ionisation energy and electronegativity increase across the period and the elements become more non-metallic in character.

The most reactive metals are found at the bottom left-hand side of the periodic table and the most reactive non-metals at the top of group 17.

Elements in group 18, the noble gases, have a very stable electronic confi guration and so are unreactive.

summary

8 Explain in you own words the: a meaning of the term electronegativity b meaning of the term core charge c relationship between electronegativity and core charge. 9 Figure 3.8 gives electronegativity values for the elements of

groups 1, 2 and 14–17 of the periodic table. a Give the symbol and name of the element that has the: i highest electronegativity ii lowest electronegativity.

b In which group do we see the: i greatest change in electronegativity as we go down the

group? ii smallest change in electronegativity as we go down the

group? c Why are the elements of group 18 usually omitted from

tables that give electronegativity values? 10 a Explain in your own words the meaning of ionisation energy. b What factors need to be considered when predicting the

trend in ionisation energies across a period?

key questions

Figure 3.13The proportion of sodium to chlorine in the compound sodium chloride (salt) always remains the same, whatever the quantity of common salt. We can demonstrate experimentally that sodium, chlorine and sodium chloride have the properties listed in Table 3.6.

1 g of sodium chloridecontains 0.393 g ofsodium and 0.607 g

of chlorine.

2 g of sodium chloridecontains 0.786 g ofsodium and 1.214 g

of chlorine.

39.3%sodium 39.3%

sodium

39.3%sodium

60.7%chlorine

60.7%chlorine

60.7%chlorine

5 g of sodium chloridecontains 1.97 g of sodium

and 3.03 g of chlorine.

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explained in detail in Chapters 5–7. For the present, we are interested in the fact that sodium chloride is just one of the large number of compounds in which a metal and a non-metal are chemically combined. Such compounds tend to have properties similar to those of sodium chloride.

chemfactMany elements form compounds with chlorine; these compounds are generally known as chlorides, for example, potassium chloride (KCl), calcium chloride (CaCl

2),

hydrogen chloride (HCl). Each compound contains a specifi c percentage of chlorine.

By contrast, water is made up of the elements hydrogen and oxygen. Both hydrogen and oxygen are gases under normal conditions, yet water is a liquid. The chemical reactivities of hydrogen and oxygen are very different from each other and from the reactivity of water. Water is not just a mixture of hydrogen and oxygen and so its properties are quite different from those of the two elements. Water is represented by the formula H2O, which means that the hydrogen and oxygen atoms are combined in the ratio of 2 : 1. We describe this unit of two hydrogen atoms combined with one oxygen atom as a water molecule. In general, a molecule is two or more non-metal atoms chemically combined. All molecules of a particular compound contain the same number of each type of atom.

Sodium (a metal) Chlorine (a non-metal) Sodium chloride

Melts at 98°C Melts at −101°C Melts at 801°C

Conducts electricity when solid

Does not conduct electricity when solid


Conducts electricity when molten (liquid)

Does not conduct electricity when liquid

Conducts electricity when molten

TABLE 3.6 Comparison of properties of sodium, chlorine and sodium chloride

Hydrogen (a non-metal) Oxygen (a non-metal) Water

Melts at −259°C Melts at −218°C Melts at 0°C







TABLE 3.7 Comparison of properties of hydrogen, oxygen and water

chemfactMany elements form compounds with oxygen; these compounds are generally known as oxides, for example, potassium oxide (K

2O), calcium oxide (CaO), sulfur

dioxide (SO2), hydrogen oxide or water

(H2O). Each compound contains a specifi c

percentage of oxygen.

As with sodium chloride, each of these properties tells us something about the particles present in hydrogen, oxygen and water and the strength of the forces between them. This will be explained in detail in Chapter 7. For the present, we are interested in the fact that water is just one of the large number of compounds in which two non-metals are chemically combined. Such compounds tend to have similar physical properties to those of water.

Just as we observe trends in the properties of the elements, so we observe similar trends in the properties of their compounds such as the chlorides and oxides.

Such observations reinforce the logic of the arrangement of elements in the periodic table. More importantly, if we come across a compound with which we are not familiar, identifying where its constituent elements are located in the periodic table gives us a good guide to its likely properties. This knowledge is also important if we wish to produce a new compound with a particular set of properties.

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chemistry in action

An American chemist, Thomas Midgley, searched for a chemical to use as a refrigerant in 1930. The chemical needed to be easily liquefi ed, odourless, non-fl ammable and non-toxic if it were to suit his purposes. By using known trends in properties of compounds of individual elements, he discovered dichlorodifl uoromethane (CF

2Cl

2), a member of a

family of compounds known as chlorofl uorocarbons (CFCs).Midgley publicly demonstrated how non-toxic the chemical was by taking a deep

breath of it and using it to blow out a candle. His discovery allowed the widespread introduction of home refrigerators and air-conditioners. CFCs have also been extensively used as the propellant in aerosol cans.

Although at the time CFCs were seen as most benefi cial because they allowed for safe and effective refrigeration, they are now being phased out due to their damaging effect on the atmosphere. CFCs are discussed further in Chapters 19 and 20.

ychemistry in action

Chemical by design

3.4

When a metal reacts with a non-metal, a compound forms that has similar properties to those of sodium chloride. It is likely to have a high melting temperature and conduct electricity in the liquid state but not as a solid.

When two or more non-metals react together, a compound forms which has physical properties similar to those of water. It is likely to have a low melting temperature and be a non-conductor of electricity.

summary

11 Elements and compounds are classifi ed as pure substances. a What is meant by the term pure substance? b In what ways are elements and compounds similar? c In what ways do elements and compounds differ?12 Sodium chloride contains 39.3% sodium. Calculate the mass of: a sodium in 15.6 g of sodium chloride b chlorine in 3.5 kg of sodium chloride c sodium chloride that contains 5 g of sodium d sodium that will be combined with 2.75 g of chlorine.13 a Refer to Table 3.6 and identify two properties that clearly

indicate that sodium chloride is chemically different from sodium and chlorine.

b Refer to Table 3.7 and identify the key property that indicates hydrogen, oxygen and water are three different chemicals.

14 Sulfur is the element below oxygen in group 16 of the periodic table. Predict the formula of the compound formed between hydrogen and sulfur.

15 Consider elements that make up the following compounds. Which of the compounds are likely to have properties similar to those of sodium chloride?

A calcium oxide (CaO) C hexane (C6H

14)

B sulfur dioxide (SO2) D potassium bromide (KBr)

key questions

Figure 3.14The use of CFCs in refrigerators and air-conditioners is being phased out because of the damage CFCs cause to the atmosphere.

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03 key terms

The periodic table

actinidesalkali metalsatomic radiuscompoundcore charged-block

electronegativityf-blockfi rst ionisation energygrouphalogenslanthanides

main group elementsmetalloidsmetalsmoleculenoble gasnon-metals

p-blockperiods-blocktransition metals

Why are element properties periodic?16 In the modern periodic (Figure 3.2), explain why: a there are two groups of elements in the s-block b there are six groups of elements in the p-block c there are ten elements in each transition series d there are fourteen elements in the actinides and the

lanthanides17 Sketch an outline of the modern periodic table, indicating the

four main blocks. Label your diagram to indicate: a the shell that is being progressively fi lled as you move

across each period b the subshell that is being fi lled in each of the main blocks18 Name an element with properties similar to those of: a carbon c iodine b rubidium d phosphorus19 In his version of the periodic table, Mendeleev arranged elements

in order of increasing atomic weight but placed tellurium (atomic weight 127.6) before iodine (atomic weight 126.9).

a Suggest why he did this? b Why is tellurium located before iodine in the modern

periodic table?

Trends in properties20 Gallium is an element in group 3, period 4 of the periodic

table. By reference to the table, predict: a whether the element is a metal or non-metal b whether a compound formed from the reaction of gallium

and oxygen will have properties similar to those of sodium chloride or those of water

21 As you move across a period, the number of subatomic particles in an atom increases but the size of an atom decreases. Why?

22 Account for the fact that it takes more energy to remove an electron from the outer shell of atoms of:

a phosphorus than magnesium b fl uorine than iodine23 Consider the elements in period 2 of the periodic table:

lithium, beryllium, boron, carbon, nitrogen, oxygen and fl uorine. Describe the changes that occur as you move across the period. Consider:

a the sizes of atoms b metallic character c electronegativity24 Chemists once called the group 18 elements the inert gases.

Why is this name less popular now?25 How do chemists explain the very low reactivity of the noble

gases?

Compounds26 Beryllium reacts with oxygen to form beryllium oxide, formula

BeO. Predict the formulas of the oxides of the other elements in group 2.

27 If calcium oxide is 71.5% calcium and 28.5% oxygen, calculate: a the mass of calcium in 4.6 g of calcium oxide b the mass of oxygen in 0.46 g of calcium oxide c the mass of oxygen you would need to react with 10 g of

calcium to form calcium oxide28 Suppose you reacted 10 g of calcium with 10 g of oxygen

until all of the calcium was converted to calcium oxide. Use your answer to Question 27c to calculate the mass of oxygen you would have left over after the reaction.

Connecting the main ideas29 Mendeleev used the properties of elements to help him form

groups in his periodic table. The following table gives some properties of elements that

have been labelled W, X, Y and Z. In each case decide if the element is most like sodium or chlorine in its reactions with other elements. (You may need to refer to Table 3.6.)

Element Melting temperature

Property

W 64°C Compound formed when W reacts with oxygen conducts electricity when molten.

X −39°C Element X conducts electricity in solid and liquid states.

Y 119°C Compound formed when Y reacts with fl uorine does not conduct electricity when molten.

Z 1540°C Compound formed when Z reacts with fl uorine conducts electricity when molten.

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The periodic table

52TTThTTTThhTTTh

52

30 What concepts or ideas are the two versions of the periodic table in Figure 3.15 trying to convey? Search the Internet for other versions of the periodic table.

1 H

2 He

3 Li

4 Be

11 Na

12 Mg

9 F

10 Ne

7 N

8 O

5 B

6 C

17 Cl

18 A

15 P

16 S

13 Al

14 Si

35 Br

36 Kr

33 As

34 Se

31 Ga

32 Ge

53 I

54 Xe

51 Sb

52 Te

49 In

50 Sn

85 At

86 Rn

83 Bi

84 Po

81 Tl

82 Pb

117 Xx

118 Xx

115 Xx

116 Xx

113 Xx

114 Xx

19 K

20 Ca

37 Rb

38 Sr

55 Cs

56 Ba

87 Fr

88 Ra

119 Xx

120 Xx

27 Co

28 Ni

25 Mn

26 Fe

23 V

24 Cr

45 Rh

46 Pd

43 Tc

44 Ru

41 Nb

42 Mo

77 Ir

78 Pt

75 Re

76 Os

73 Ta

74 W

109 Mt

110 Xx

107 Bh

108 Hs

105 Db

106 Sg

21 Sc

22 Ti

39 Y

40 Zr

71 Lu

72 Hf

103 Lr

104 Rf

29 Cu

30 En

47 Ag

48 Cd

79 Au

80 Hg

111 Xx

112 Xx

69 Tm

70 Yb

67 Ho

68 Er

101 Md

102 No

99 Es

100 Fm

63 Eu

64 Gd

61 Pm

62 Sm

59 Pr

60 Nd

95 Am

96 Cm

93 Np

94 Pu

91 Pa

92 U

57 La

58 Co

89 Ac

90 Th

65 Tb

66 Dy

97 Bk

98 Cf

167 168 165 166 163 164 169 170 159 160 157 158 155 156 153 154 161 162 151 152 149 150 145 146 143 144 141 142 139 138 140 147 148

Rn

Xe

Kr

Ar

Ne

He F Cl

Br I

At

Po Te Se S O

N C

Al

Ga

In

Tl Hg

Cd

Zn

112

111

Au

Ag

Cu

Ni

Co

Rh

Ir

Mt

Hs

Hs

Ru

Fe

Mn Te Re Ns

Sg W

Mo

Cr

V

Nb Ta Ha

Rf Hf

Lu

Yb Tm

Er Ho

Dr

Tb Gd Eu

Sm Pm

Nd Pr

Ce La

Ac

Ra

Ba

Sr

Ca

Mg Be

H

B

Li

Na

K

Rb

Cs

Fr

Th Pa

U Np

Pu Am

Cm

Bk

Cf

Es

Fm Md

No Lr

Zr

Ti

Sc

Y

Pd

Pt

110

Si

Ge

Sn

Pb

P

As

Sb

Bi

TRANSITION METALS

LANTHANIDES AND ACTINIDES

SUPE

RACTI

NIDES

Figure 3.15a The periodic spiral of Professor Thoedor Benfey.b Albert Tarantola’s orbital periodic table.

a

b

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chapter 03 the periodic 0 table

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