05 periodic trends note - chemwithbilal.com

Periodic Trends

9 The Periodic Table: chemical periodicity

This topic illustrates the regular patterns in some physical properties of the elements in the Periodic Table.

9.1 Periodicity of physical properties of the elements in the third period

Bilal Hameed 1 Periodic Trends

Periodic Trends

Periodic Trends 2 Bilal Hameed

Cambridge International AS and A Level Chemistry 9701 syllabus Syllabus content

29Back to contents page www.cie.org.uk/alevel

Inorganic chemistry

9 The Periodic Table: chemical periodicity

This topic illustrates the regular patterns in chemical and physical properties of the elements in the Periodic Table.

Learning outcomesCandidates should be able to:

9.1 Periodicity of physical properties of the elements in the third period

a) describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements (see the Data Booklet )

b) explain qualitatively the variation in atomic radius and ionic radius

c) interpret the variation in melting point and electrical conductivity in terms of the presence of simple molecular, giant molecular or metallic bonding in the elements

d) explain the variation in first ionisation energy (see the Data Booklet)

e) explain the strength, high melting point and electrical insulating properties of ceramics in terms of their giant structure; to include magnesium oxide, aluminium oxide and silicon dioxide

9.2 Periodicity of chemical properties of the elements in the third period

a) describe the reactions, if any, of the elements with oxygen (to give Na2O, MgO, Al 2O3, P4O10, SO2, SO3), chlorine (to give NaCl , MgCl 2, Al 2Cl 6, SiCl 4, PCl 5) and water (Na and Mg only)

b) state and explain the variation in oxidation number of the oxides (sodium to sulfur only) and chlorides (sodium to phosphorus only) in terms of their valence shell electrons

c) describe the reactions of the oxides with water

(treatment of peroxides and superoxides is not required)

d) describe and explain the acid/base behaviour of oxides and hydroxides including, where relevant, amphoteric behaviour in reaction with acids and bases (sodium hydroxide only)

e) describe and explain the reactions of the chlorides with water

f) interpret the variations and trends in 9.2(b), (c), (d) and (e) in terms of bonding and electronegativity

g) suggest the types of chemical bonding present in chlorides and oxides from observations of their chemical and physical properties

9.3 Chemical periodicity of other elements

a) predict the characteristic properties of an element in a given Group by using knowledge of chemical periodicity

b) deduce the nature, possible position in the Periodic Table and identity of unknown elements from given information about physical and chemical properties

1

PERIODIC TRENDSElements show trends in their physical and chemical properties across periods and down groups.

The periodic table is divided into groups, periods and blocks. The International Union of Pure and Applied Chemistry, IUPAC, now recommends that the groups in the periodic table should be numbered 1 to 18. • Groups 1 and 2 remain the same as before – classified as the s-block. • The transition elements now become Group 3 to 12 – classified as the d-block. • Groups 3 to 7 now become Groups 13 to 17 and the noble gases become Group

18 – classified as the p-block.

All the elements in a group contain the same number of electrons in their outer shell. For example, all the elements in group 1 have one electron in their outer shell, which is in an s-orbital. All the elements in group 17 have seven electrons in their outer orbit, arranged s2p5.

The horizontal rows are called periods. All elements in the same period have the same number of shells containing electrons. For example, all the elements in the third period, Na to Ar, have electrons in the first, second and third shells.

INORGANIC CHEMISTRY

188

in rows or periods across the table. The elements in a period have different physical and chemical properties, but trends become apparent in these properties as we move across a period.

The s, p, d and f blocksThe fi rst two groups form the s block (Groups 1 and 2). These have the outer electronic confi gurations ns1 and ns2, respectively, where n is the number of the shell. The last six groups form the p block (Groups 13 to 18), in which the p sub-shell is being progressively fi lled. These elements have the outer electronic confi gurations ns2 np1 to ns2 np6.

In between the s and p blocks is the d block (Groups 3–12), in which the d sub-shell is being progressively fi lled. Our study of the d block is largely restricted to the elements scandium to zinc, and these elements have the outer electronic confi gurations 4s2 3d1 to 4s2 3d10, as we shall see in Topic 24.

About a quarter of the known elements belong to the f block. The fi rst row, the 4f, includes all the elements from cerium (Ce) to lutetium (Lu) inclusive. In this block, the f sub-shell in the fourth principal shell is being progressively fi lled. In the past, these elements were called the ‘rare earths’. This was a misnomer, as their abundance in the Earth’s crust is fairly large – for some it is comparable to that of lead. However, it is true that they are mostly quite thinly spread, so mining them is quite expensive. They are now called the lanthanoids, as their properties are similar to those of the element lanthanum (La) that precedes them in the Periodic Table. Although they are not particularly rare, the lanthanoids are diffi cult to purify from one another because their chemical properties are nearly identical. The elements of the second row of the f block, the 5f, are called the actinoids. All the actinoids are radioactive, and most have to be made artifi cially.

10.3 Periodic trends in the elements of the third period (sodium to argon)In this section we shall look at the trends in the properties of the elements and their compounds in the third period of the Periodic Table.

AppearanceThe elements on the left of the Periodic Table have low values of ionisation energy and electronegativity, and so they show the properties associated with metallic bonding (see section 4.11), for example, they are shiny and conduct electricity. In the middle of the Periodic Table, elements with higher values of ionisation energy and electronegativity are semiconductors: they have a dull shine to them and are poor conductors of electricity (typically 10−12 times that of a metal). The elements on the right of the Periodic Table, with the highest values of ionisation energy and electronegativity, are dull in appearance and are such poor conductors that they are used as electrical insulators (their conductivities are only about 10−18 times that of a metal).

10_02 Cam/Chem AS&A2

Barking Dog Art

f block

d block

p block1

2

3

4

5

6

7

Li Be

Na Mg

K Ca

Rb Sr

Cs Ba

Fr

Sc

Y

Lu

Lr

Ti

Zr

Hf

Rf

V

Nb

Ta

Db

Cr

Mo

W

Sg

Mn

Tc

Re

Bh

Fe

Ru

Os

Hs

Co

Rh

Pr

Mt

Ni

Pd

Pt

Cu

H

Ag

Au

Zn

Cd

Hg

Ga

Pn

B

Al

TI

Ge

Sn

C

Si

Pb

As

Sb

N

P

Bi

Se

Te

O

S

Po

Br

P

F

Cl

At

Kr

Xe

Ne

He

Ar

Rn

Ds Rg Cn Fl113 115 Lv 117 118Ra

La

Ac

Ce

Th

Pr

Pa

Nd

U

Pm

Np

Sm

Pu

Eu

Am

Gd

Cm

Tb

Bk

Dy

Cf

Ho

Es

Er

Fm

Tm

Md

Yb

No

1 Group

Peri

od

s block2 13 14 15 16 17

43 5 6 7 8 9 10 11 12

18

118

Figure 10.2 A modern form of the Periodic Table. The elements shown in red are good conductors of electricity, and the ones shown in blue are poor conductors. The pink shading indicates that the principal form of the element is a semiconductor. As only a few atoms of the elements at the end of the Periodic Table have been made, these have been left white.

181333_10_AS_Chem_BP_186-201.indd 188 18/09/14 1:49 PM

116

Cha

pter

8 T

he p

erio

dic

tabl

e an

d pe

riod

icit

y

4

Be

20

Ca

38

1

H1.0

3

Li6.9 9.0

Sr

87

Fr88

Ra

56

Ba

19

K39.1 40.1 45.0 47.9 50.9 52.0

24.3

11

23.0

Na12

Mg

21

Sc22

Ti23

V24

Cr

41

Nb42

Mo

105

Db106

Sg

25

54.9

Mn26

55.8

Fe27

58.9

102.9 106.4

58.7

Co28

Ni29

63.5

Cu

47

107.9Ag

111

Rg

79

Au

109

Mt110

Ds

77

Ir78

195.1192.2190.2186.2183.8180.9178.5137.3

55

Cs132.9

37

Rb85.5 87.6 88.9 91.2 92.9 95.9

Pt

45

Rh46

Pd

107

Bh108

Hs

75

Re76

Os

43

Tc44

101.1Ru

73

Ta74

W

89–103 104

Rf

57-7172

Hf

39

Y40

Zr

1

2

30

65.4

Zn

5

10.8

B6

12.0

C7

14.0

N8

16.0

O9

19.0

F10

20.2

Ne

85

At86

Rn

53

126.9I

54

131.3Xe

35

79.9

Br36

83.8

Kr

17

35.5

Cl18

39.9

Ar

2

4.0

He

83

209.0207.2204.4200.6197.0

Bi84

Po

51

121.8Sb

52

127.6Te

33

74.9

As34

79.0

Se

15

31.0

P16

32.1

S

81

Tl82

Pb

49

114.8In

50

118.7Sn

31

69.7

Ga32

72.6

Ge

13

27.0

Al14

28.1

Si

48

112.4Cd

80

Hg

Cn Fl Lv112 114 116

1176543 8 12

18Keyatomic numbersymbolrelative atomic mass

89

Ac

57

138.9

La58

140.1

Ce59

Pr60

144.2

Nd61

144.9

Pm

96

Cm

64

157.2

Gd

94

Pu95

Am

62

150.4

Sm63

152.0

Eu

92

238.1

U93

Np

140.9

90

232.0

Th91

Pa102

No103

Lr

70

Yb71

Lu

100

Fm101

Md

68

167.3

Er69

173.0 175.0168.9

Tm

98

Cf99

Es

66

162.5

Dy67

164.9

Ho

97

Bk

65

158.9

Tb

109

1413 15 16 17

Figure 8.1 The periodic table.

The International Union of Pure and Applied Chemistry, IUPAC, now recommends that the groups in the periodic table should be numbered 1 to 18.

! Groups 1 and 2 remain the same as before – classified as the s-block.

! The transition elements now become Group 3 to 12 – classified as the d-block.

! Groups 3 to 7 now become Groups 13 to 17 and the noble gases become Group 18 – classified as the p-block.

The blocks are shown in Figure 8.2.

p- blockblock

s-

blockd-

Figure 8.2 Blocks of the periodic table.

The vertical columns are called groups and the horizontal rows are called periods. Trends in the groups (vertical columns) and periods (horizontal rows) reflect the structures of the atoms of the elements within them, and these in turn affect the chemical properties of the elements. These repeating properties demonstrate periodicity.

Key term

Periodicity is a repeating pattern, in either physical or chemical properties, across different periods.

9781471827068_OCR_A_Level_Chemistry.indb 116 30/03/15 2:44 PM

1

Bilal Hameed MarginalizerBilal Hameed 3 Periodic Trends

1

PERIODIC TRENDSThe table is also divided into blocks:

The s-block consists of the elements in groups 1 and 2. An s-block element has its highest-energy electron in an s-orbital, i.e. the last electron added goes into an s-orbital.

The elements in the s-block are reactive metals. These metals (including potassium, sodium, calcium and magnesium) are all high in the activity (electrochemical) series. They have lower densities, lower melting points and lower boiling points than most other metals and they form stable, involatile ionic compounds.

Group 1: Also called the Alkali Metals. Soft, highly reactive group of metals. ns1.

Group 2. Also called the Alkaline Earth Metals. Fairly reactive metals. ns2

The d-block elements occupy space in the periodic table between Group 2 and Group 13. The d-block contains the elements scandium to zinc in period 4 and those elements below them.

A d-block element cannot be defined in terms of energy because the energy level of the d-orbitals is altered by the presence of electrons in the outer s-orbital. It can only be defined in aufbau terms, i.e. the last electron added to a d-block element goes into a d-orbital.

These metals (including chromium, iron, copper, zinc and silver) are much less reactive than the metals in Groups 1 and 2.

116

Cha

pter

8 T

he p

erio

dic

tabl

e an

d pe

riod

icit

y

4

Be

20

Ca

38

1

H1.0

3

Li6.9 9.0

Sr

87

Fr88

Ra

56

Ba

19

K39.1 40.1 45.0 47.9 50.9 52.0

24.3

11

23.0

Na12

Mg

21

Sc22

Ti23

V24

Cr

41

Nb42

Mo

105

Db106

Sg

25

54.9

Mn26

55.8

Fe27

58.9

102.9 106.4

58.7

Co28

Ni29

63.5

Cu

47

107.9Ag

111

Rg

79

Au

109

Mt110

Ds

77

Ir78

195.1192.2190.2186.2183.8180.9178.5137.3

55

Cs132.9

37

Rb85.5 87.6 88.9 91.2 92.9 95.9

Pt

45

Rh46

Pd

107

Bh108

Hs

75

Re76

Os

43

Tc44

101.1Ru

73

Ta74

W

89–103 104

Rf

57-7172

Hf

39

Y40

Zr

1

2

30

65.4

Zn

5

10.8

B6

12.0

C7

14.0

N8

16.0

O9

19.0

F10

20.2

Ne

85

At86

Rn

53

126.9I

54

131.3Xe

35

79.9

Br36

83.8

Kr

17

35.5

Cl18

39.9

Ar

2

4.0

He

83

209.0207.2204.4200.6197.0

Bi84

Po

51

121.8Sb

52

127.6Te

33

74.9

As34

79.0

Se

15

31.0

P16

32.1

S

81

Tl82

Pb

49

114.8In

50

118.7Sn

31

69.7

Ga32

72.6

Ge

13

27.0

Al14

28.1

Si

48

112.4Cd

80

Hg

Cn Fl Lv112 114 116

1176543 8 12

18Keyatomic numbersymbolrelative atomic mass

89

Ac

57

138.9

La58

140.1

Ce59

Pr60

144.2

Nd61

144.9

Pm

96

Cm

64

157.2

Gd

94

Pu95

Am

62

150.4

Sm63

152.0

Eu

92

238.1

U93

Np

140.9

90

232.0

Th91

Pa102

No103

Lr

70

Yb71

Lu

100

Fm101

Md

68

167.3

Er69

173.0 175.0168.9

Tm

98

Cf99

Es

66

162.5

Dy67

164.9

Ho

97

Bk

65

158.9

Tb

109

1413 15 16 17

Figure 8.1 The periodic table.

The International Union of Pure and Applied Chemistry, IUPAC, now recommends that the groups in the periodic table should be numbered 1 to 18.

! Groups 1 and 2 remain the same as before – classified as the s-block.

! The transition elements now become Group 3 to 12 – classified as the d-block.

! Groups 3 to 7 now become Groups 13 to 17 and the noble gases become Group 18 – classified as the p-block.

The blocks are shown in Figure 8.2.

p- blockblock

s-

blockd-

Figure 8.2 Blocks of the periodic table.

The vertical columns are called groups and the horizontal rows are called periods. Trends in the groups (vertical columns) and periods (horizontal rows) reflect the structures of the atoms of the elements within them, and these in turn affect the chemical properties of the elements. These repeating properties demonstrate periodicity.

Key term

Periodicity is a repeating pattern, in either physical or chemical properties, across different periods.


2

Marginalizer Bilal HameedPeriodic Trends 4 Bilal Hameed

1

PERIODIC TRENDSThe p-block contains the elements in groups 13 to 17 and group 18 (the noble gases). A p-block element has its highest-energy electron in a p-orbital, i.e. the last electron added goes into a p-orbital.

Some of the elements in the p-block are metals, like tin, lead and bismuth. They are usually low in the activity (electrochemical) series and they have some resemblances to non-metals. Other elements in the p-block are non-mental.

Group 17. Also called the Halogens. Very reactive group of non-metals. ns2 np5

Group 18. Also called the Noble Gases because they do not react. Form few compounds. ns2 np6

182

The periodic table and periodicity

In this section you will learn to:

• Observe the area where metals and non-metals separate

• Define a ‘metalloid’

• Compare atomic electrical conductivity of the metals, metalloids and non-metals

9.4 Metals, non-metals and metalloids

Although the periodic table does not classify elements as metals and non-metals, there is a fairly obvious division between the two (‘fairly obvious’, but, not ‘clear cut’). Separating the elements into either metals or non-metals is rather like trying to separate all the shades of grey into either black or white. The fairly obvious division between metals and non-metals is shown by a thick stepped line in Figure 9.5. The 20 or so non-metals are packed into the top right-hand corner above the thick stepped line. Some of the elements next to the thick steps, such as germanium, arsenic and antimony, have similarities to both metals and non-metals and it is difficult to place these in one class or the other. Chemists sometimes use the name metalloid for these elements which are difficult to classify one way or the other.

Figure 9.8 shows a classification of elements as metals, metalloids and non-metals on the basis of their electrical conductivity. In this classification:

B C N O F

Al Si P S Cl

Ga Ge As Se Br

In Sn Sb Te I

Tl Pb

Li Be

Na Mg

K Ca

Rb Sr

Cs BaFr Ra

Bi Po At

Ne

Ar

Kr

Xe

Rn

Metals

Metals

MetalloidsNon-metals

Non-metals

d-block andf-block metals

Fig 9.8 Classifying the elements as metals, metalloids and non-metals on the basis of their electrical conductivity. In this figure, metalloids are coloured grey

NOTE*The atomic electrical conductivity (atomic conductance) is the conductivity of a block of the substance 1 cm2 in cross-section but long enough to contain one mole of atoms of the element. It is a measure of the conductivity of one mole of atoms of the element.

DEFINITIONA metalloid is an element with some properties like metals and other properties like non-metals

1 Metals are good conductors of electricity with atomic electrical conductivity* greater than 10−3 ohm−1 cm−4.

2 Metalloids are poor conductors of electricity with atomic electrical conductivity usually less than 10−3 but greater than 10−5 ohm−1 cm−4.

3 Non-metals are virtually non-conductors (insulators). Their atomic electrical conductivity is usually less than 10−10 ohm−1 cm−4.

Notice in Figure 9.8 that the cell for carbon is shaded less heavily than those for the other metalloids. This is because carbon exists as three different allotropes – graphite, a poor conductor classed as a metalloid, plus diamond and fullerenes, insulators classed as non-metals.

In spite of problems such as this, the classification of elements into metals, metalloids and non-metals is useful and convenient.

147

Learning objectives:➔ State the location of the s-,

p-, and d-blocks of elements in the Periodic Table.

Speci!cation reference: 3.2.1

The Periodic Table is a list of all the elements in order of increasing atomic number. You can predict the properties of an element from its position in the table. You can use it to explain the similarities of certain elements and the trends in their properties, in terms of their electronic arrangements.

The structure of the Periodic TableThe Periodic Table has been written in many forms including pyramids and spirals. The one shown below is one common layout. Some areas of the Periodic Table are given names. These are shown in Figure 1.

1

2

3

4

5

6

7

1 2 3 4 5 6 7 0

metalloids

alka

li m

etal

s

halo

gens

nobl

e (r

are

or in

ert)

gase

s

transition metals

alka

line

earth

met

als

Lanthanides

Actinides

lanthanides

actinides

▲ Figure 1 Named areas of the Periodic Table

Metals and non-metalsThe red stepped line in Figure 1 (the ‘staircase line’) divides metals (on its left) from non-metals (on its right). Elements that touch this line, such as silicon, have a combination of metallic and non-metallic properties. They are called metalloids or semi-metals. Silicon, for example, is a non-metal but it looks quite shiny and conducts electricity, although not as well as a metal.

8.1 The Periodic Table

+ History of the Periodic TableThe development of the Periodic Table is one of the greatest achievements in chemistry. Credit for the !nal version goes !rmly to a Russian, Dmitri Mendeleev, in 1869. He realised that there were undiscovered elements. He left spaces for some unknown elements, and arranged the known elements so that similar elements lined up in columns. Since then, new elements have been discovered that !t into the gaps he left. Mendeleev even accurately

predicted the properties of the missing elements, con!rming the success of his Periodic Table.

Many other scientists contributed to the Periodic Table. Research on the internet the parts played by Jons Jacob Berzelius, Robert Bunsen and Gustav Kircho", Alexandre Béguyer de Chancourtois, Marie Curie, Sir Humphry Davy, Julius Lothar Meyer, Henry Moseley, John Newlands, Sir William Ramsay, and Glenn T Seaborg.

8 Periodicity

3


1

PERIODIC TRENDS

The periodicity of atomic properties In this section we ll look at the trends in atomic and ionic radii, first ionisation energies and electronegativity as you go across a period.

To understand how such atomic properties vary we have to understand all the forces at play inside an atom.

3.1 Periodic table 87

■ Electron arrangement and the periodic tableIt is the electrons in the outer or valence shell that determine the chemical and physical properties of the chemical element. The position of a chemical element in the periodic table is related to its electron arrangement. The period number indicates the number of shells in the atom of the element. All chemical elements in the same period have the same number of shells. In groups 1 and 2, the number of valence electrons is equal to the group number. In groups 13 to 18, the number of valence electrons is equal to the group number minus 10.

Figure 3.6 shows how the electron arrangement of a chemical element is related to its group and period number. This so-called ‘short form’ of the periodic table omits the transition elements.

■ Figure 3.2 Samples of the period 3 elements: sodium, magnesium, aluminium, silicon, phosphorus, sulfur, chlorine and argon

■ Figure 3.3 From left to right: thiocyanate ions, iron(III) ions and complex ions formed by the reaction between iron(III) and thiocyanate ions

■ Figure 3.4 A sample of hydrated nickel chloride, NiCl2.6H2O; nickel is a transition metal

■ Figure 3.5 British £2 coin showing an outer gold-coloured nickel-brass ring made from 76% copper, 20% zinc and 4% nickel and an inner silver-coloured cupro-nickel disc made from 75% copper and 25% nickel

H

K Ca

Na Mg

Li Be

Si P S Cl Ar

C N O F Ne

He

B

Al

1

1

2

3

4

2 13 14 15 16 17 18

1

Group

Period

19 20

11 12 13 14 15 16 17 18

3 4 5 6 7 8 9 10

2

1

2,8,8,1 2,8,8,2

2,8,1 2,8,2

2,1 2,2

2,8,3 2,8,4 2,8,5 2,8,6 2,8,7 2,8,8

2,4 2,5 2,6 2,7 2,8

2

2,3

■ Figure 3.6 The short form of the periodic table, showing the first 20 chemical elements and their electron arrangements

Based on the electron arrangements of the elements (Chapter 12), the periodic table can be divided into four blocks of elements (Figure 3.7):■ s-block elements■ p-block elements■ d-block elements■ f-block elements.

829055_03_IB_Chemistry_085-113.indd 87 18/05/15 9:28 am185

The intermolecular forces between simple molecules in these elements involve only weak induced dipole attractions (Section 4.15). So, their melting points and boiling points are low.

Electrical conductivity

The periodicity of electrical conductivity for the elements in Periods 2 and 3, like the periodicity in melting points and boiling points, depends critically on their structure and bonding.

Metals, such as sodium, magnesium and aluminium, on the left of each period with mobile, delocalised electrons in their structure are good conductors of electricity. When a battery is attached to them, electrons flow out of the metals into the positive terminal of the battery. At the same time, electrons flow into the metal from the negative terminal of the battery.

On the other hand, metalloids and non-metals with molecular structures have no mobile electrons. This means that their electrical conductivity is almost nil.

7 a Which elements occur at or near the peaks on the graph in Figure 9.9? What type of structure do these elements have?

b Suggest a reason why the electrical conductivity rises from Na → Mg → Al.

Q U E S T I O N

Table 9.3 shows the electronic (shell and sub-shell) structures for the elements in Periods 2 and 3.

Elements in the same group of the periodic table have similar electron configurations. For example, Group I elements (Li to Fr) have one electron in their outer shells (ns1). Group II elements (Be to Ra) have two electrons in their outer shells (ns2) and Group VII elements (F to At) have seven outer shell electrons (ns2np5). As these electron configurations recur in a periodic pattern, it would not be surprising to find that atomic properties, such as atomic radii, ionic radii and ionisation energies, show similar periodicity.

Atomic radii

The atomic radii of atoms can be obtained from X-ray analysis and electron density maps. Using these techniques, it is possible to measure the distance between the nuclei of atoms and then estimate the radius of individual atoms.

The atomic radii of metals are obtained by measuring the distance between the nuclei of neighbouring atoms in metal crystals (Figure 9.10 (a)). The atomic radius is simply half of the inter-nuclear distance.

9.7 The periodicity of atomic properties

Period 2 Li Be B C N O F Ne

Electron shell structure 2, 1 2, 2 2, 3 2, 4 2, 5 2, 6 2, 7 2, 8

Electron sub-shell structure 1s22s1 1s22s2 1s22s22p1 1s22s22p2 1s22s22p3 1s22s22p4 1s22s22p5 1s22s22p6

Period 3 Na Mg Al Si P S Cl Ar

Electron shell structure 2, 8, 1 2, 8, 2 2, 8, 3 2, 8, 4 2, 8, 5 2, 8, 6 2, 8, 7 2, 8, 8

Electron sub-shell structure 1s22s22p6

3s1

1s22s22p6

3s2

1s22s22p6

3s23p1

1s22s22p6

3s23p2

1s22s22p6

3s23p3

1s22s22p6

3s23p4

1s22s22p6

3s23p5

1s22s22p6

3s23p6

Table 9.3 The electronic structures of the elements in Periods 2 and 3

KEY POINTElements in the same group of the periodic table have similar electron configurations. This results in similar atomic properties.

9.7 The periodicity of atomic properties

In this section you will learn to:

• Give the electronic structure of Period 2 and 3 elements and compare the number of electrons in the last sub-shell

• Describe and explain changes in atomic and ionic radii as you go across a period

• Describe and explain other trends in atomic properties such as first ionisation energies and electronegativity

150

Learning objectives:➔ Describe the trends in melting

and boiling temperatures of the elements in Period 3.

➔ Explain these trends in terms of bonding and structure.


The Periodic Table reveals patterns in the properties of elements. For example, every time you go across a period you go from metals on the left to non-metals on the right. This is an example of periodicity. The word periodic means recurring regularly.

Periodicity and properties of elements in Period 3Periodicity is explained by the electron arrangements of the elements.

• The elements in Groups 1, 2, and 3 (sodium, magnesium, and aluminium) are metals. They have giant structures. They lose their outer electrons to form ionic compounds.

• Silicon in Group 4 has four electrons in its outer shell with which it forms four covalent bonds. The element has some metallic properties and is classed as a semi-metal.

• The elements in Groups 5, 6, and 7 (phosphorus, sulfur, and chlorine) are non-metals. They either accept electrons to form ionic compounds, or share their outer electrons to form covalent compounds.

• Argon in Group 0 is a noble gas – it has a full outer shell and is unreactive.

Table 1 shows some trends across Period 3 (see Figure 1). Similar trends are found in other periods.

1234567

1 2 3 4 5 6 7 0

Na Mg Al Si P S Cl Ar

▲ Figure 1 The Periodic Table with Period 3 highlighted

Study tip• Remember that when a

molecular substance melts, the covalent bonds remain intact but the van der Waals forces break.

• Learn the formulae P4, S8, Cl2.

Group 1 2 3 4 5 6 7 0Element sodium magnesium aluminium silicon phosphorus sulfur chlorine argonElectron

arrangement [Ne] 3s1 [Ne] 3s2 [Ne] 3s2 3p1 [Ne] 3s2 3p2 [Ne] 3s2 3p3 [Ne] 3s2 3p4 [Ne] 3s2 3p5 [Ne] 3s2 3p6

s-block p-blockmetals semi-metal non-metals noble gas

Structure of element giant metallic

macromolecular (giant covalent)

molecular atomicP4 S8 Cl2 Ar

Melting point, Tm / K 371 922 933 1683

317 (white)

392 (monoclinic)

172 84

Boiling point, Tb / K 1156 1380 2740 2628

553 (white)

718 238 87

▼ Table 1 Some trends across Period 3

8.2 Trends in the properties of elements of Period 3

4


1

PERIODIC TRENDSForces that act on electrons in an atom The forces that act on electrons in an atom are electromagnetic. They obey Coulomb’s law, which states that the magnitude of the force between two charged objects is directly proportional to the product of their charges and inversely proportional to the square of the distance between their centres.

It simplifies to:

• opposite charges attract; like charges repel • the bigger the charges, the stronger is the force • the further apart the particles, the weaker is the force

This means that electrons are attracted towards the nucleus. The greater the atomic number, the stronger is the force of attraction. Electrons that are further away from the nucleus are attracted less than those closer to the nucleus.

In addition, because they have the same charge, electrons repel each other. Because they are more densely packed, the inner electrons repel outer electrons much more than the other outer electrons repel each other.

27Forces that act on electrons in an atom

The periodic table can be used to predict the electron con! guration of an element because it indicates the order in which orbitals are ! lled (Figure 2.17).

[He]

Core

2s

1s

[Ne] 3s[Ar] 4s 3d

4d[Kr] 5s

2p3p4p5p

[Xe] 6s

Figure 2.17 The periodic table and the order of ! lling orbitals

Forces that act on electrons in an atomThe forces that act on electrons in an atom are electromagnetic. They obey Coulomb’s law, which states that the magnitude of the force between two charged objects is directly proportional to the product of their charges and inversely proportional to the square of the distance between their centres. This is expressed mathematically as:

force ∝ q+q−

r2

where q+ and q− are the charges on the two objects and r is the distance between their centres.

The rules are:

● opposite charges attract; like charges repel● the bigger the charges, the stronger is the force● the further apart the particles, the weaker is the force

This means that electrons are attracted towards the nucleus (Figure 2.18). The greater the atomic number, the stronger is the force of attraction. Electrons that are further away from the nucleus are attracted less than those closer to the nucleus.


Shielding and the effective nuclear chargeEffective nuclear charge across period 1In a hydrogen atom, the nucleus has a charge of +1 and there is one electron in a 1s-orbital. There are no forces of repulsion, so the electron feels the full force of attraction of a +1 charge.

In a helium atom, the nucleus has a charge of +2 and there are two electrons in the 1s-orbital. These electrons repel each other slightly. The result is that the net force of attraction between the nucleus and each electron is slightly less than that between a +2 charge and one electron. Therefore, the helium nucleus is said to have an e! ective nuclear charge of slightly less than 2.

Effective nuclear charge across period 2The situation becomes more complicated for lithium and the remaining elements. In lithium, the nucleus has a charge of +3 and the outer 2s electron is strongly repelled

Key termThe effective nuclear charge is the net charge on the nucleus, after allowing for the electrons in orbit around the nucleus shielding its full charge.

+ –

–

–

Attraction

Repulsion

Repulsion

Figure 2.18 The forces acting on an electron in an atom

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27Forces that act on electrons in an atom

The periodic table can be used to predict the electron con! guration of an element because it indicates the order in which orbitals are ! lled (Figure 2.17).

[He]

Core

2s

1s

[Ne] 3s[Ar] 4s 3d

4d[Kr] 5s

2p3p4p5p

[Xe] 6s

Figure 2.17 The periodic table and the order of ! lling orbitals

Forces that act on electrons in an atomThe forces that act on electrons in an atom are electromagnetic. They obey Coulomb’s law, which states that the magnitude of the force between two charged objects is directly proportional to the product of their charges and inversely proportional to the square of the distance between their centres. This is expressed mathematically as:

force ∝ q+q−

r2

where q+ and q− are the charges on the two objects and r is the distance between their centres.

The rules are:

● opposite charges attract; like charges repel● the bigger the charges, the stronger is the force● the further apart the particles, the weaker is the force

This means that electrons are attracted towards the nucleus (Figure 2.18). The greater the atomic number, the stronger is the force of attraction. Electrons that are further away from the nucleus are attracted less than those closer to the nucleus.


Shielding and the effective nuclear chargeEffective nuclear charge across period 1In a hydrogen atom, the nucleus has a charge of +1 and there is one electron in a 1s-orbital. There are no forces of repulsion, so the electron feels the full force of attraction of a +1 charge.

In a helium atom, the nucleus has a charge of +2 and there are two electrons in the 1s-orbital. These electrons repel each other slightly. The result is that the net force of attraction between the nucleus and each electron is slightly less than that between a +2 charge and one electron. Therefore, the helium nucleus is said to have an e! ective nuclear charge of slightly less than 2.

Effective nuclear charge across period 2The situation becomes more complicated for lithium and the remaining elements. In lithium, the nucleus has a charge of +3 and the outer 2s electron is strongly repelled

Key termThe effective nuclear charge is the net charge on the nucleus, after allowing for the electrons in orbit around the nucleus shielding its full charge.

+ –

–

–

Attraction

Repulsion

Repulsion

Figure 2.18 The forces acting on an electron in an atom


5


1

PERIODIC TRENDSFactors that effect the force of attraction on outer electrons Nuclear charge When the nuclear charge becomes more positive (due to the presence of additional protons), its attraction on all the electrons increases.

Atomic Radius As the distance of the outer electrons from the nucleus increases, the attraction of the positive nucleus for the negatively charged electrons falls.

Shielding Effect The outer or valence electrons are repelled by all the other electrons in the atom in addition to being attracted by the positively charged nucleus. The outer electrons are shielded from the attraction of the nucleus by the shielding effect.

2 Atomic structure74

The third shell can hold a maximum of 18 electrons. However, when there are eight electrons in the third shell there is a degree of stability and the next two electrons enter the fourth shell. For the transition metals beyond calcium the additional electrons enter the third shell until it contains the maximum of 18 electrons. In addition, the second and subsequent shells are divided into a number of sub-shells. Atoms (other than hydrogen) also rearrange their electrons before they can form chemical bonds with other atoms.

This process is called hybridization (Chapter 14). An important concept introduced in Chapter 3 and also in Chapter 12 is electron shielding (Figure 2.50). The electrons in the different shells experience different attractive forces due to the presence of other electrons. The outer electrons experience the most shielding.

+

–

–

–

––

–

––

–

–

These electrons shield the outerelectron from the nucleus

This electron does notfeel the full effect of the positivecharge of the nucleus–

+

–

–

–

–

–

––

–

–

■ Figure 2.50 Electron shielding

ToK LinkJacob Bronowski: ‘One aim of the physical sciences has been to give an exact picture of the material world. One achievement… has been to prove that this aim is unattainable’. What are the implications of this claim for the aspirations of natural sciences in particular and for knowledge in general?

This claim is probably related to ‘modern’ (approximately 100 years old) physics such as Einstein’s theory of special relativity and quantum physics, including the Uncertainty Principle. An exact description (picture) of the material or physical world is impossible. Quantum mechanics means that the world is ‘fuzzy’ at the atomic and molecular level, and there are limits of experimental precision dictated by the Uncertainty Principle. Matter on the atomic level is ‘schizophrenic’ due to its wave–particle duality. The material world is described by a series of scientific models, all of which are limited and incomplete as descriptions of physical phenomena.

There is no absolute knowledge in science and the scientific method assumes that there is a material world of objects and phenomena existing out there that is independent of the observers, the scientists. However, some physicists might question the assertion that the material world is independent of the observers. Thus, Schrödinger’s cat is both alive and dead until the observation is made, i.e. the box is opened, the wave function is collapsed, and one of the eventualities – alive or dead – is manifested.

Niels Bohr wrote, ‘It is wrong to think that the task of physics is to find out how nature is. Physics concerns what we can say about nature.’

ToK LinkHeisenberg’s Uncertainty Principle states that there is a theoretical limit to the precision with which we can know the momentum and the position of a particle. What are the implications of this for the limits of human knowledge?

In 1927 the German physicist Heisenberg stated the Uncertainty Principle, which is the consequence of the dual behaviour of matter and radiation (de Broglie’s hypothesis). It states that it is impossible to determine simultaneously the exact position and exact momentum (or velocity) of an electron (or any sub-atomic particle) along a given direction.

Mathematically, it can be described by the equation:

Δx × Δp ≥ h

4π

where Δx is the uncertainty in position and 6p is the uncertainty in momentum (or velocity) of the particle. The momentum of the electron is the product of its mass (m) and velocity (v). If the position of the electron is known with a high degree of accuracy (6x is small), then the velocity of the electron will be uncertain. However, if the velocity of the electron is known precisely, then the position of the electron will be uncertain (6x will be large).

For an electron whose mass is 9.110 × 10 −31 kg,

∆v × ∆x ≥ h

4π m = 6.63 × 10–34 J s

4 × 3.1416 × 9.110 × 10–31 kg ≈ 10−4 m2 s−1

10 Find out about the thought experiment ‘Schrödinger’s Cat’

829055_02_IB_Chemistry_052-084.indd 74 18/05/15 9:26 am

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Explaining periodic patterns

Atomic radius

Atomic radii decrease

Atomic radii increase

Figure 8.3 Periodicity of atomic radii in the periodic table.

Across a period, the atomic radius decreases from left to right (Figure 8.3).

From one atom of an element to the next across a period:

! the charge of the nucleus increases

! the shielding remains the same

Across the period the effect of the nuclear charge on the outer electrons increases and the atomic radii decreases.

Down a group, the atomic radius increases.

From one atom of an element to the next down a group:


! the shielding increases

Down the group the effect of the nuclear charge on the outer electrons decreases and the atomic radii increases.

outer electronshielded by 2inner electrons



Li Na K

Figure 8.5 Atoms of the Group 1 elements lithium, sodium and potassium.

Test yourself

1 Put the following elements in order of increasing atomic radius. Justify your answers:a) Mg, S, Sib) Mg, K, Alc) Si, Cl, K

Na

shieldingelectrons infull shells

10–

11+

1–

Figure 8.4 Shielding effect of inner shell electrons reduces the pull of the nucleus in the outer shell.



Q Electron configurationIonization energyThe first ionization energy is the minimum energy per mole required to remove electrons from one mole of isolated gaseous atoms to form one mole of gaseous unipositive ions under standard thermodynamic conditions. For example, the first ionization energy of chlorine is the energy required to bring about the reaction:

Cl(g) → Cl+(g) + e−

The electron is removed from the outer sub-shell (energy sub-level) of the chlorine atom (that is, a 3p electron). Table 12.1 gives some examples of ionizations, and in each case the ionization energy, which is the enthalpy change for the equation. Ionization energies are listed Table 8 of the IB Chemistry data booklet.

Element Ionization equation First ionization energy/kJ mol−1

Oxygen O(g) → O+(g) + e− 1314Sulfur S(g) → S+(g) + e− 1000Copper Cu(g) → Cu+(g) + e− 745

Factors that affect ionization energyValues of ionization energies depend on the following factors:Q the size of the atom (or ion)Q the nuclear chargeQ the shielding effect.

Atomic radiusAs the distance of the outer electrons from the nucleus increases, the attraction of the positive nucleus for the negatively charged electrons falls. This causes the ionization energy to decrease. Hence, ionization energy decreases as the atomic or ionic radius increases.

Nuclear chargeWhen the nuclear charge becomes more positive (due to the presence of additional protons), its attraction on all the electrons increases. This causes the ionization energy to increase.

Shielding effectThe outer or valence electrons are repelled by all the other electrons in the atom in addition to being attracted by the positively charged nucleus. The outer electrons are shielded from the attraction of the nucleus by the shielding effect (an effect of electron–electron repulsion) (Figure 12.13).

Q Table 12.1 Selected ionization energies

7 Calculate the wavelength (in m) of electromagnetic radiation with a frequency of 1368 kHz. Deduce which part of the electromagnetic spectrum it belongs to.

8 Calculate the frequency of yellow light with a wavelength of 5800 × 10−8 cm.

9 The laser used to read information from a compact disc has a wavelength of 780 nm. Calculate the energy associated with one photon of this radiation.

Q Figure 12.13 Electrostatic forces operating on the outer or valence electron in a lithium atom

3+

–

––

nuclearpull

repulsion frominner shell of

electrons (’shielding’)

E = hcλ

= (6.63 × 10−34 J s × 3.00 × 108 m s−1)

600 × 10−9 m = 3.31 × 10−19 J per photon

4.0 × 10−17 J

3.31 × 10−19 J/photon = 1.2 × 102 photons

Calculate the number of photons with wavelength 4.00 nm that can provide 1 joule of energy.

For one photon: E = hcλ

= (6.63 × 10−34 J s × 3.00 × 108 m s−1)

4 × 10−9 m = 4.9725 × 10−17 J

So for one joule: 1

4.9725 × 10−17 = 2.01 × 1016 photons


6


1

PERIODIC TRENDSShielding and the effective nuclear charge The effective nuclear charge is the net charge on the nucleus, after allowing for the electrons in orbit around the nucleus shielding its full charge.

Effective nuclear charge across period 1 In a hydrogen atom, the nucleus has a charge of +1 and there is one electron in a 1s-orbital. There are no forces of repulsion, so the electron feels the full force of attraction of a +1 charge.

In a helium atom, the nucleus has a charge of +2 and there are two electrons in the 1s-orbital. These electrons repel each other slightly. The result is that the net force of attraction between the nucleus and each electron is slightly less than that between a +2 charge and one electron. Therefore, the helium nucleus is said to have an effective nuclear charge of slightly less than 2.

Effective nuclear charge across period 2

The situation becomes more complicated for lithium and the remaining elements. In lithium, the nucleus has a charge of +3 and the outer 2s electron is strongly repelled by the two inner 1s electrons. The nucleus is shielded by the inner electrons and the effective nuclear charge is approximately +1. This is the +3 nuclear charge, minus the effect of two negatively charged screening electrons.

The next element is beryllium. The nucleus has a charge of +4, there are two 1s electrons shielding the nucleus and two 2s electrons that also repel each other slightly. Therefore, the effective nuclear charge is not exactly +2 (+4 nuclear charge minus the effect of the two negative inner electrons) — it is slightly less than +2 because of the extra repulsion by the two electrons in the outer orbit.

The situation is slightly more complicated with the next element, boron. The atomic number of boron is five (a nuclear charge of +5). There are two 1s electrons that shield the outer electrons from the nucleus. The two 2s electrons are closer to the nucleus than the single 2p electron and they repel it. Therefore, the effective nuclear charge is significantly less than the +3 value predicted by the simplified idea that effective nuclear charge is equal to the atomic number of the element minus the number of inner-shell electrons.








H

K Ca

Na Mg

Li Be

Si P S Cl Ar

C N O F Ne

He

B

Al

1

1

2

3

4

2 13 14 15 16 17 18

1

Group

Period

19 20

11 12 13 14 15 16 17 18

3 4 5 6 7 8 9 10

2

1

2,8,8,1 2,8,8,2

2,8,1 2,8,2

2,1 2,2

2,8,3 2,8,4 2,8,5 2,8,6 2,8,7 2,8,8

2,4 2,5 2,6 2,7 2,8

2

2,3




7


1

PERIODIC TRENDS

Similar arguments apply to other periods — the effective nuclear charge increases across a period, but does not increase by as much as +1 between successive elements.

12.1 Electrons in atoms 443

2s 2p2p

In general, the shielding effect is most effective if the electrons are close to the nucleus. Consequently, electrons in the first shell (energy level), where there is high electron density, have a stronger shielding effect than electrons in the second shell, which in turn have a stronger shielding effect than electrons in the third shell. Electrons in the same shell exert a relatively small shielding effect on each other.

Figure 12.14 shows the first ionization energies for the chemical elements of periods 1, 2 and 3. The general increase in ionization energy across each period is due to the increase in nuclear charge. This occurs because across the period each chemical element has one additional proton, which increases the nuclear charge by +1.

Q Figure 12.14 First ionization energies for periods 1, 2 and 3

2500

2000

1500

1000

500

01 105 15 20

Atomic number, Z

He

H

Ne

Ar

Li NaK

Firs

t ion

izat

ion

ener

gy/k

J m

ol–1 shielding

force

Li Be B C N O F Ne

electrostaticattraction

towardspositivenucleus

Q Figure 12.15 A diagram illustrating how the balance between shielding and nuclear charge changes across period 2

boron atom, B1s 2s 2p

boron ion, B+

1s 2s 2p

beryllium atom, Be1s 2s 2p

beryllium ion, Be+

1s 2s 2p

Q Figure 12.16 Orbital notations for boron and beryllium atoms and their unipositive ions

Q Figure 12.17 Electron density clouds of the 2s and 2p orbitals (only one lobe shown). The dotted line shows the extent of the 1s orbital; the 2s electron can partially penetrate the 1s orbital, increasing its stability

The increase in nuclear charge increases the force of attraction on all the electrons, so they are held closer and hence more strongly. Each additional electron across a period enters the same shell (energy level) and hence the increase in shielding is minimal (Figure 12.15).

Although the general trend is for the ionization energy to increase across the period, there are two distinct dips in ionization energy across periods 2 and 3 (Figure 12.14). These dips can only be explained using an orbital model of electronic structure.

The first decrease in each period is the result of a change in the sub-shell (sub-level) from which the electron is lost and a change in electron shielding. These have a greater effect than the increase in nuclear charge and decrease in atomic radius. In period 2, this first decrease occurs between the elements beryllium and boron. When it is ionized, the beryllium atom (1s22s2) loses a 2s electron, whereas a boron atom (1s22s22p1) loses a 2p electron (Figure 12.16). More energy is required to remove an electron from the lower energy 2s orbital in beryllium than from the higher energy 2p orbital in boron. Although the 2s and 2p sub-levels are in the same shell, the energy difference is relatively large. Recall (Chapter 2) that the energy gap between shells and sub-levels becomes smaller with an increase in shell number. In addition, a single electron in the 2p sub-level is more effectively shielded by the inner electrons than the 2s2 electrons (Figure 12.17).









H

K Ca

Na Mg

Li Be

Si P S Cl Ar

C N O F Ne

He

B

Al

1

1

2

3

4

2 13 14 15 16 17 18

1

Group

Period

19 20

11 12 13 14 15 16 17 18

3 4 5 6 7 8 9 10

2

1

2,8,8,1 2,8,8,2

2,8,1 2,8,2

2,1 2,2

2,8,3 2,8,4 2,8,5 2,8,6 2,8,7 2,8,8

2,4 2,5 2,6 2,7 2,8

2

2,3




86

The core charge increases across the period. This means that the outer-shell electrons of chlorine therefore experience a greater attraction to the nucleus than does the outer-shell electron of sodium.

Atomic radius decreases from left to right across period 3 due to the increasing attraction experienced by the outer-shell electrons. These outer-shell electrons are all in the third electron shell of the atoms; however, as the core charge increases, the electrostatic attraction between the outer-shell electrons and the nucleus increases. This has the effect of pulling the electrons in closer to the nucleus and making the atom smaller.

The trend in ionic radii is not as clear as for atomic radii (see fi gure 3.2.6). For the metals (sodium to aluminium) in period 3, the ionic radius decreases across the period. Silicon can be represented as a positive (Si4+) or negative (Si4−) ion. For the non-metals, the ionic radius decreases from the phosphorus (P3−) to the chloride ion (Cl−) (fi gure 3.2.7). A positive ion has a smaller ionic radius than the original atom, due to the loss of the valence electrons, and a negative ion has a larger ionic radius than the original atom, since the addition of extra negative charges introduces more electron–electron repulsion. Negative ions have a larger radius than positive ions, as they have one more shell of

electrons than the positive ions.

The increase in fi rst ionization energy and electronegativity from left to right across period 3 can also be explained by the increasing core charge. As the core charge increases, it becomes increasingly diffi cult to remove an electron from the outer shell of the atom (fi rst ionization energy). Similarly, the increasing electrostatic attraction of outer-shell electrons for the nucleus results in a greater power of attraction for electrons in the outer shell (electronegativity).

1 State and explain how:a the fi rst ionization energy of strontium compares with that

of magnesiumb the electronegativity of selenium compares with that of oxygen.

2 a Explain what is meant by the term core charge.b How is core charge used to explain the trend in atomic radius of the

period 3 elements?

3 Explain why the electronegativity of fl uorine is higher than that of magnesium.

4 a Compare the atomic radii of magnesium and chlorine.b Explain the difference you have described.

5 a Compare the fi rst ionization energy of phosphorus and chlorine. b Explain the difference you have described.

6 a State the trend in the ionic radii from Na+ to Al3+.b State the trend in the ionic radii from Si4− to Cl−.c Explain the trend in part b.

11+

innerelectronsshield thevalenceelectronfrom thenucleus

electronattractedby aneffectivecharge of +1

this electronexperiencesa strongerattractionthan theelectron in sodium

17+

sodiumcore charge = 11 – 10 = +1

chlorinecore charge = 17 – 10 = +7

Figure 3.2.8 Core charge can be used to explain trends within the periodic table.

• atomic radius decreases• electronegativity increases• first ionization energy increases• ionic radius decreases

Figure 3.2.9 Trends in properties across a period.

Section 3.2 Exercises

Summary of periodic trends

WORKSHEET 3.3 Periodic table trends

86

The core charge increases across the period. This means that the outer-shell electrons of chlorine therefore experience a greater attraction to the nucleus than does the outer-shell electron of sodium.

Atomic radius decreases from left to right across period 3 due to the increasing attraction experienced by the outer-shell electrons. These outer-shell electrons are all in the third electron shell of the atoms; however, as the core charge increases, the electrostatic attraction between the outer-shell electrons and the nucleus increases. This has the effect of pulling the electrons in closer to the nucleus and making the atom smaller.

The trend in ionic radii is not as clear as for atomic radii (see fi gure 3.2.6). For the metals (sodium to aluminium) in period 3, the ionic radius decreases across the period. Silicon can be represented as a positive (Si4+) or negative (Si4−) ion. For the non-metals, the ionic radius decreases from the phosphorus (P3−) to the chloride ion (Cl−) (fi gure 3.2.7). A positive ion has a smaller ionic radius than the original atom, due to the loss of the valence electrons, and a negative ion has a larger ionic radius than the original atom, since the addition of extra negative charges introduces more electron–electron repulsion. Negative ions have a larger radius than positive ions, as they have one more shell of

electrons than the positive ions.

The increase in fi rst ionization energy and electronegativity from left to right across period 3 can also be explained by the increasing core charge. As the core charge increases, it becomes increasingly diffi cult to remove an electron from the outer shell of the atom (fi rst ionization energy). Similarly, the increasing electrostatic attraction of outer-shell electrons for the nucleus results in a greater power of attraction for electrons in the outer shell (electronegativity).

1 State and explain how:a the fi rst ionization energy of strontium compares with that

of magnesiumb the electronegativity of selenium compares with that of oxygen.

2 a Explain what is meant by the term core charge.b How is core charge used to explain the trend in atomic radius of the

period 3 elements?

3 Explain why the electronegativity of fl uorine is higher than that of magnesium.

4 a Compare the atomic radii of magnesium and chlorine.b Explain the difference you have described.

5 a Compare the fi rst ionization energy of phosphorus and chlorine. b Explain the difference you have described.

6 a State the trend in the ionic radii from Na+ to Al3+.b State the trend in the ionic radii from Si4− to Cl−.c Explain the trend in part b.

11+

innerelectronsshield thevalenceelectronfrom thenucleus

electronattractedby aneffectivecharge of +1

this electronexperiencesa strongerattractionthan theelectron in sodium

17+

sodiumcore charge = 11 – 10 = +1

chlorinecore charge = 17 – 10 = +7

Figure 3.2.8 Core charge can be used to explain trends within the periodic table.

• atomic radius decreases• electronegativity increases• first ionization energy increases• ionic radius decreases

Figure 3.2.9 Trends in properties across a period.

Section 3.2 Exercises

Summary of periodic trends

WORKSHEET 3.3 Periodic table trends

&)-

6idb^X�gVY^jhThe atomic radius is basically used to describe the size of an atom. The larger the atomic radius, the larger the atom.

The atomic radius is usually taken to be half the internuclear distance in the element. For example, in a diatomic molecule such as chlorine, where two identical atoms are joined together, the atomic radius would be de! ned as shown in Figure 4.8.

Atomic radius increases down a group.

This is because, as we go down a group in the periodic table the atoms have increasingly more electron shells. For example, potassium has four shells of electrons but lithium has only two:

electron is more di" cult to remove from an argon atom. The argon atom is also smaller than the sodium atom and, therefore, the outer electron is closer to the nucleus and more strongly held.

&& &-

CV 6g

Vidb^X�gVY^jh

;^\jgZ�)#-� I]Z�Vidb^X�gVY^jh�d[�X]adg^cZ�Vidbh�^c�V�bdaZXjaZ# @

A^

@

A^Atomic radius decreases across a period.

Although the nuclear charge is higher for K, the number of electrons and hence the repulsion between electrons is also greater, and this counteracts any e# ects due to a greater number of protons in the nucleus.

Figure 4.9 shows the variation of the atomic radius across period 3 in the periodic table.

:miZch^dc

It is possible to de! ne two di# erent atomic radii: the covalent radius and the van der Waals’ radius.

8


1

PERIODIC TRENDSEffective nuclear charge in a group The nuclear charge and the number of inner shielding electrons increase by the same amount in a group. This leads to an assumption that the effective nuclear charge acting on the outer electrons of the elements in the same group of the periodic table hardly varies, but this is a simplification of a more complex situation.

For instance, sodium has 11 protons and, therefore, a nuclear charge of +11. It has two electrons in the first shell and eight in the second shell. These ten electrons shield the outer, third-shell electron very efficiently and the effective nuclear charge is close to +1 (+11 − 2 − 8 = +1).

Potassium has 19 protons and, therefore, a nuclear charge of +19. The outer, fourth-shell electron is shielded from the nucleus by two electrons in the first shell, eight electrons in the second shell and eight electrons in the third shell, making a total of 18 inner shielding electrons. From sodium to potassium, the number of protons has increased by eight and the number of shielding electrons has also increased by eight. Therefore, potassium also has a similar effective nuclear charge close to +1.

118

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icit

yExplaining periodic patterns

Atomic radius



Figure 8.3 Periodicity of atomic radii in the periodic table.

Across a period, the atomic radius decreases from left to right (Figure 8.3).

From one atom of an element to the next across a period:


! the shielding remains the same


Down a group, the atomic radius increases.

From one atom of an element to the next down a group:


! the shielding increases





Li Na K

Figure 8.5 Atoms of the Group 1 elements lithium, sodium and potassium.

Test yourself

1 Put the following elements in order of increasing atomic radius. Justify your answers:a) Mg, S, Sib) Mg, K, Alc) Si, Cl, K

Na

shieldingelectrons infull shells

10–

11+

1–

Figure 8.4 Shielding effect of inner shell electrons reduces the pull of the nucleus in the outer shell.


9


1

PERIODIC TRENDSAtomic radius The radius of an atom is found by measuring the distance between the nuclei of two touching atoms, and then halving that distance.

In general the atomic radius of an atom is determined by the balance between two opposing factors:

• the shielding effect by the electrons of the inner shell(s) – this makes the atomic radius larger. The shielding effect is the result of repulsion between the electrons in the inner shell and those in the outer or valence shell.

• the nuclear charge (due to the protons) – this is an attractive force that pulls all the electrons closer to the nucleus. With an increase in nuclear charge, the atomic radius becomes smaller.

Atomic radius decreases across a period Remember that effective nuclear charge increases across the period. This means that the outer-shell electrons of chlorine therefore experience a greater attraction to the nucleus than does the outer-shell electron of sodium.

Atomic radius decreases from left to right across period 3 due to the increasing attraction experienced by the outer-shell electrons. These outer-shell electrons are all in the third electron shell of the atoms; however, as the effective nuclear charge increases, the electrostatic attraction between the outer-shell electrons and the nucleus increases. This has the effect of pulling these outer-shell electrons in closer to the nucleus and making the atom smaller.

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The reason that atomic radius decreases across a period is basically the same reason electronegativity and ionisation energy increase: an increase in nuclear charge across the period but no signi! cant increase in shielding.

Sodium and chlorine have the same number of inner shells of electrons (and hence the amount of shielding is similar); however, chlorine has a nuclear charge of 17+ whereas sodium has a nuclear charge of only 11+. This means that the outer electrons are pulled in more strongly in chlorine than in sodium and the atomic radius is smaller.

>dc^X�gVY^jhThe ionic radius is a measure of the size of an ion.

In general, the ionic radii of positive ions are smaller than their atomic radii, and the ionic radii of negative ions are greater than their atomic radii.

For instance, Figure 4.10 (overleaf ) shows a comparison of the atomic and ionic radii (1+ ion) for the alkali metals. Each 1+ ion is smaller than the atom from which it is formed (by loss of an electron).

Na is larger than Na+ as it has one extra shell of electrons – the electronic con! guration of Na is 2, 8, 1, whereas that of Na+ is 2, 8. Also, they both have the same nuclear charge pulling in the electrons (11+), but there is a greater amount of electron–electron repulsion in Na, as there are 11 electrons compared with only 10 in Na+. The electron cloud is therefore larger in Na than in Na+, as there are more electrons repelling for the same nuclear charge pulling the electrons in.

The fact that negative ions are larger than their parent atoms can be seen by comparing the sizes of halogen atoms with their ions (1−) in Figure 4.11 (overleaf ). Cl− is larger than Cl, because it has more electrons for the same nuclear charge and, therefore, greater repulsion between electrons. Cl has 17 electrons and 17 protons in the nucleus. Cl− also has 17 protons in the nucleus, but it has 18 electrons. The repulsion between 18 electrons is greater than between 17 electrons, so the electron cloud expands as an extra electron is added to a Cl atom to make Cl−.

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152

Learning objectives:➔ Describe the trends in atomic

radius and !rst ionisation energy of the elements in Period 3.

➔ Explain these trends.Speci!cation reference: 3.2.1

Some key properties of atoms, such as size and ionisation energy, are periodic, that is, there are similar trends as you go across each period in the Periodic Table.

Atomic radiiThese tell us about the sizes of atoms. You cannot measure the radius of an isolated atom because there is no clear point at which the electron cloud density around it drops to zero. Instead half the distance between the centres of a pair of atoms is used, see Figure 1.

The atomic radius of an element can differ as it is a general term. It depends on the type of bond that it is forming – covalent, ionic, metallic, van der Waals, and so on. The covalent radius is most commonly used as a measure of the size of the atom. Figure 2 shows a plot of covalent radius against atomic number.

(Even metals can form covalent molecules such as Na2 in the gas phase. Since noble gases do not bond covalently with one another, they do not have covalent radii and so they are often left out of comparisons of atomic sizes.)

The graph shows that:

• atomic radius is a periodic property because it decreases across each period and there is a jump when starting the next period

• atoms get larger down any group.

Why the radii of atoms decrease across a periodYou can explain this trend by looking at the electronic structures of the elements in a period, for example, sodium to chlorine in Period 3, as shown in Figure 3.

As you move from sodium to chlorine you are adding protons to the nucleus and electrons to the outer main shell, which is the third shell. The charge on the nucleus increases from +11 to +17. This increased charge pulls the electrons in closer to the nucleus. There are no additional electron shells to provide more shielding. So the size of the atom decreases as you go across the period.

▲ Figure 1 Atomic radii are taken to be half the distance between the centres of a pair of atoms

r

▲ Figure 2 The periodicity of covalent radii. The noble gases are not included because they do not form covalent bonds with one another

0 5 2010 15

Li Na

Period 2Period 3

K

atomic number

covalentradius / nm

0.00

0.24

0.06

0.12

0.18

Hint1 nm is 1 × 10–9 m

Study tipIt is a common mistake to think that atoms increase in size as you cross a period. While the nuclei have more protons (and neutrons) the radius of the atom depends on the size of the electron shells.

atom

size ofatom

atomic(covalent)radius / nm

nuclearcharge

2,8,2

Na Mg Al

0.156 0.136 0.125

11+ 12+ 13+

2,8,1 2,8,3

Si

0.117

14+

0.110

15+

2,8,5

0.104

16+

2,8,6

0.099

17+

2,8,7

P S Cl

2,8,5 2,8,6 2,8,72,8,4

▲ Figure 3 The sizes and electronic structures of the elements sodium to chlorine

More trends in the properties of the elements of Period 3

8.310


1

PERIODIC TRENDS

Summary: Across a period, the atomic radius decreases from left to right. From one atom of an element to the next across a period: • the charge of the nucleus increases • the shielding remains the same


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2 Atomic structure and the periodic table (Topic 1)30

Atomic radiusThe radius of an atom is found by measuring the distance between the nuclei of two touching atoms, and then halving that distance.

Trends across a periodEven though extra electrons are being added, the atoms get smaller going across a period from left to right. From lithium to !uorine, the outer electrons are all in the 2nd shell, being screened by the 1s2 electrons. The increasing number of protons in the nucleus as you go across the period pulls the electrons in more tightly. This is slightly o"set by the increased repulsion of the electrons from each other. However, the net e"ect is a decrease in radius (Figure 2.19). This pattern is repeated in the 3rd period.

50Li Be B C N O F Na Mg Al Si P S Cl K Ca

Ato

mic

rad

ius/

pm

100

150

200

250

2nd period 3rd period

Figure 2.19 Trends in atomic radius

Trends in a groupDown a group, the trend is for atomic radii to increase steadily because of the increase in the number of occupied shells:

● Lithium: electron structure 2,1 — two occupied shells.● Sodium: electron structure 2,8,1 — three occupied shells.● Potassium: electron structure 2,8,8,1 — four occupied shells.

Size of positive ionsA positive ion is always smaller than its neutral atom. If an atom loses all its outer electrons, the radius of the resulting ion is much smaller than the atomic radius. This is because:

● there is one fewer shell of electrons● there are fewer electrons in the positive ion than in the atom, so the electron–

electron repulsion is less, causing a further reduction in the radius

For ions with the same electron con#guration (e.g. Na+, Mg2+ and Al3+), the ion with the greatest charge will have the smallest radius.

The three positive ions in Table 2.6 have the same number of electrons (ten) arranged in the same way (1s2 2s2 2p6). Therefore, the electron–electron repulsion is the same.


11


1

PERIODIC TRENDSAtomic radius increases down a group Down a group from top to bottom, atomic radii increase. In each new period the outer-shell electrons enter a new energy level so are located further away from the nucleus. This has a greater effect than the increasing nuclear charge because of shielding by the inner-shell electrons (effective nuclear charge stays nearly the same.)

Summary: Down a group, the atomic radius increases. From one atom of an element to the next down a group:

• the charge of the nucleus increases • the shielding effect increases (effective nuclear charge stays the same)

• increase in the number of shells


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The variation of ionic radius across a period is not a clear-cut trend, as the type of ion changes from one side to the other. Thus positive ions are formed on the left-hand side of the period and negative ions on the right-hand side.

For positive ions there is a decrease in ionic radius as the charge on the ion increases, but for negative ions the size increases as the charge increases (Figure 4.12).

Let us consider Na+ and Mg2+: both ions have the same electronic con! guration, but Mg2+ has one more proton in the nucleus (Figure 4.13). Because there is the same number of electrons in both ions, the amount of electron–electron repulsion is the same; however, the higher nuclear charge in Mg2+ means that the electrons are pulled in more strongly and so the ionic radius is smaller.

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3.2 Periodic trends 97

Moving down a group, both the nuclear charge and the shielding effect increase. However, the outer electrons enter new shells. So, although the nucleus gains protons, the electrons are not only further away, but also more effectively screened by an additional shell of electrons (Figure 3.22).

atomic radii decrease

atomic radii increaseatomic radii increase

■ Figure 3.22 Summary of trends in periodicity in atomic radii in the periodic table

Atomic number0 10

F

ClBr

I At

20 30 40 50 7060 90800

300

200

100Ato

mic

radi

us/p

m

■ Figure 3.21 Bar chart showing the variation of atomic radii in group 17

■ Figure 3.20 Bar chart showing the variation of atomic radii in group 1

Atomic number0 10

Li

Na

KRb

Cs Fr

20 30 40 50 7060 90800

300

200

100Ato

mic

radi

us/p

m

■ Table 3.4 The variation of atomic radii in group 1

Atom Atomic number Atomic radius/pm

Li 3 152

Na 11 186

K 19 231

Rb 37 244

Cs 55 262

Fr 87 270

Atom Atomic number Atomic radius/pm

F 9 58

Cl 17 99

Br 35 114

I 53 133

At 85 140

■ Table 3.5 The variation of atomic radii in group 17

Ionic radii for ions of the same charge also increase down a group for the same reason (Tables 3.6 and 3.7). Ionic radii are the radii for ions in a crystalline ionic compound (Figure 3.23).


12


1

PERIODIC TRENDSHow is atomic radius measured We know that electrons in atoms are in located in atomic orbitals, which are regions of space where there is a high probability of finding an electron. This means that the position of the electron is not fixed, so we cannot measure the radius of the atom in the same way as we measure the radius of a circle. One way of overcoming this problem and finding the radius of an atom is to measure the distance between the nuclei of two closest atoms of the element and dividing it by 2.

The atomic radii of atoms can be obtained from X-ray analysis and electron density maps. Using these techniques, it is possible to measure the distance between the nuclei of atoms and then estimate the radius of individual atoms.

The atomic radius of an element can differ as it is a general term. It depends on the type of bond that it is forming – covalent, ionic, metallic, van der Waals, and so on.

The atomic radii of metals are obtained by measuring the distance between the nuclei of neighbouring atoms in metal crystals. The atomic radius is simply half of the inter-nuclear distance. This is sometimes referred to as metallic radius.

The atomic radii of non-metals are obtained from the distance between the nuclei of similar atoms joined by a covalent bond. So, for non-metals, the atomic radius is half of the covalent bond length. Because of this link to covalent bonds, the atomic radii of non-metals are sometimes called covalent radii.

152

Learning objectives:➔ Describe the trends in atomic

radius and !rst ionisation energy of the elements in Period 3.

➔ Explain these trends.Speci!cation reference: 3.2.1

Some key properties of atoms, such as size and ionisation energy, are periodic, that is, there are similar trends as you go across each period in the Periodic Table.

Atomic radiiThese tell us about the sizes of atoms. You cannot measure the radius of an isolated atom because there is no clear point at which the electron cloud density around it drops to zero. Instead half the distance between the centres of a pair of atoms is used, see Figure 1.

The atomic radius of an element can differ as it is a general term. It depends on the type of bond that it is forming – covalent, ionic, metallic, van der Waals, and so on. The covalent radius is most commonly used as a measure of the size of the atom. Figure 2 shows a plot of covalent radius against atomic number.

(Even metals can form covalent molecules such as Na2 in the gas phase. Since noble gases do not bond covalently with one another, they do not have covalent radii and so they are often left out of comparisons of atomic sizes.)

The graph shows that:

• atomic radius is a periodic property because it decreases across each period and there is a jump when starting the next period

• atoms get larger down any group.

Why the radii of atoms decrease across a periodYou can explain this trend by looking at the electronic structures of the elements in a period, for example, sodium to chlorine in Period 3, as shown in Figure 3.

As you move from sodium to chlorine you are adding protons to the nucleus and electrons to the outer main shell, which is the third shell. The charge on the nucleus increases from +11 to +17. This increased charge pulls the electrons in closer to the nucleus. There are no additional electron shells to provide more shielding. So the size of the atom decreases as you go across the period.

▲ Figure 1 Atomic radii are taken to be half the distance between the centres of a pair of atoms

r

▲ Figure 2 The periodicity of covalent radii. The noble gases are not included because they do not form covalent bonds with one another

0 5 2010 15

Li Na

Period 2Period 3

K

atomic number

covalentradius / nm

0.00

0.24

0.06

0.12

0.18

Hint1 nm is 1 × 10–9 m

Study tipIt is a common mistake to think that atoms increase in size as you cross a period. While the nuclei have more protons (and neutrons) the radius of the atom depends on the size of the electron shells.

atom

size ofatom

atomic(covalent)radius / nm

nuclearcharge

2,8,2

Na Mg Al

0.156 0.136 0.125

11+ 12+ 13+

2,8,1 2,8,3

Si

0.117

14+

0.110

15+

2,8,5

0.104

16+

2,8,6

0.099

17+

2,8,7

P S Cl

2,8,5 2,8,6 2,8,72,8,4

▲ Figure 3 The sizes and electronic structures of the elements sodium to chlorine

More trends in the properties of the elements of Period 3

8.3

Atomic radiiAtomic radii measure the size of atoms in crystals and molecules. Chemists useX-ray diffraction and other techniques to measure the distance between thenuclei of atoms. The atomic radius of an atom cannot be defined preciselybecause it depends on the type of bonding and on the number of bonds.

Atomic radii for metals are calculated from the distances between atoms inmetal crystals (metallic radii). The atomic radii for non-metals are calculatedfrom the lengths of covalent bonds in crystals or molecules (covalent radii).

Atomic radii decrease from left to right across a period. Across the period Nato Ar, atomic radii fall from 0.191 nm for sodium to 0.099 nm for chlorine. Fromone element to the next across a period the charge on the nucleus increases byone as the number of electrons in the same outer shell increases by one.Shielding by electrons in the same shell is limited, so the ‘effective nuclearcharge’ increases and the electrons are drawn more tightly to the nucleus.

Periodic properties

99

13 Use data on the CD-ROM to explore whether there is any patternin the boiling points of the chlorides of elements in Periods 2 and 3.

14 Why do you think the first ionisation energies of elements decreasewith atomic number in every group of the periodic table?

Figure 7.7 !Periodicity in the first ionisation energiesof the elements.

2500

2000

1500

1000

500

01 5 10 15 20

Firs

tion

isat

ion

energ

y/kJ

mol

–1

Atomic number

He

Ne

Argroup 0

group 2group 1

H

N

B

CO

F

Al

Si

SP

Cl

BeMg

Ca

Li NaK

Figure 7.8 "Atomic radii and the internuclear distancein a molecule and a crystal.

atomic radius

atomic radius

non-metallic molecule

metal crystal

Data



15 Arrange these elements in order ofatomic radius: Al, B, C, K and Na.

16 Which atom or ion in each of thesepairs has the larger radius?a) Cl or Cl–

b) Al or N

Test yourself

Figure 7.9 "Periodicity of atomic radii in the periodic table.

13


1

PERIODIC TRENDSHow is atomic radius measured In the simple molecular structures of non-metals, one molecule touches the next and it is sometimes useful to compare the distances between neighbouring atoms which are not chemically bonded. This distance in non-metal crystals is called the van der Waals radius.

Consider a group of gaseous argon atoms. When two argon atoms collide with one another there is very little penetration of their electron cloud densities. Argon does not form a diatomic species. If argon is frozen in the solid phase the atoms would touch each other but would not be chemically bonded. In this case the distance between the argon atoms could be measured and hence non-bonding atomic radius. The non-bonding atomic radius is often termed the van der Waals’ radius.

atoms could be measured and hence Rnb

could be found (!gure 2). The non-bonding atomic radius is often termed the van der Waals’ radius.

Section 9 of the Data booklet provides data for the covalent atomic radii of the elements. The general term “atomic radius” is used to represent the mean bonding atomic radius obtained from experimental data over a wide range of elements and compounds. Note that the bonding atomic radius is always smaller than the non-bonding atomic radius. The approximate bond length between two elements can also be estimated from their atomic radii.

For example, for the interhalogen compound BrF:

atomic radius of bromine = 117 pm

atomic radius of "uorine = 60 pm

bond length of Br-F = 177 pm

Compare this with the experimental bond length of Br-F in the gas phase (176 nm).

Quick questionPredict the bond lengths in:

a) iodine monobromide, IBr

b) trichloromethane (chloroform), CHCl3.

E!ective nuclear charge and screening e!ectIn an atom the negatively charged electrons are attracted to the positively charged nucleus. A valence or outer-shell electron is also repelled by the other electrons in the atom. The core electrons in the inner non-valence energy levels of the atom reduce the positive nuclear charge experienced by a valence electron. This e!ect of reducing the nuclear charge experienced by an electron is termed screening or shielding.

The net charge experienced by an electron is termed the effective nuclear charge, Zeff. This is the nuclear charge, Z, (representing the atomic number) minus

the charge, S, that is shielded or screened by the core electrons:

Ze! = Z - S

where Z = actual nuclear charge (atomic number) and S = screening or shielding constant.

Ze! can be worked out using Slater’s rules. You can read about these rules in advanced textbooks on inorganic chemistry, but you are not required to calculate Ze! using Slater’s rules as part of the IB Chemistry Diploma programme. You do need to understand the principle of screening and for our purposes you can consider S as a parameter related to the number of core electrons in an atom.

▲ Figure 2 Atoms of argon in the solid phase. The atoms are touching but not chemically bonded. The non-bonding atomic radius of argon Rnb is 188 pm (d = 376 pm)

d = 2Rnb

77

3 . 2 P E R I O D I C T R E N D S

14


1

PERIODIC TRENDSThe Data booklet provides data for the atomic radii of the elements. The general term “atomic radius” is used to represent the mean bonding atomic radius obtained from experimental data. Note that the bonding atomic radius (metallic and covalent) is always smaller than the non-bonding atomic radius (van der Waals).

INORGANIC CHEMISTRY

190

Phosphorus, sulfur and chlorine all form small covalent molecules, P4, S8 and Cl2 respectively. When these substances melt, it is only necessary to break weak intermolecular bonds and not strong interatomic attractions. The melting points decrease in the order sulfur > phosphorus > chlorine (see Figure 10.5), the intermolecular bonds becoming weaker as the molecules become smaller.

The melting point of argon is low, as the attraction between the argon atoms is very small.


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0

0.10

0.20

proton number

11 12 13 14 15 16 17 18

ato

mic

rad

ius/

nm

Na

Mg

Al

Si

Cl

Ar

PS

Period 3

Figure 10.4 Atomic radius and proton number

Electrical conductivitySodium, magnesium and aluminium are metals. They have delocalised electrons that are free to move in the lattice of cations (see section 4.11). Silicon is a semiconductor. The other elements in the third period form covalent bonds with no free electrons, and so are insulators with almost no electrical conductivity.

Figure 10.5 Variation of melting point with proton number


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Li

Be

C

B

N O F

Ne

Na

MgAl

Si

P S

ClAr

K

Ca

ScTi

V Cr

Mn

Fe Co Ni

Cu

Zn

Ga

Ge

As

SeBr

Kr

3500

3000

2500

2000

1500

1000

500

0

proton number1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36

mel

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oin

t/°C

181333_10_AS_Chem_BP_186-201.indd 190 18/09/14 1:49 PM

3 Periodicity96

3.2 Periodic trends – elements show trends in their physical and chemical properties across periods and down groups

First ionization energyThe first ionization energy is the minimum energy required to remove one mole of electrons from one mole of gaseous atoms (under standard thermodynamic conditions of 25 °C and 1 atm).

In general: X(g) → X+(g) + e−

For example, the first ionization energy of hydrogen is given by the following equation:

H(g) → H+(g) + e− ∆H = +1310 kJ mol−1

The amount of energy required to carry out this process for a mole of hydrogen atoms is 1310 kilojoules.

Atoms of each element have different values of first ionization energy.

ElectronegativityThe electronegativity of an atom is the ability or power of an atom in a covalent bond to attract shared pairs of electrons to itself. The greater the electronegativity of an atom, the greater its ability to attract shared pairs of electrons to itself.

Electronegativity values are usually based on the Pauling scale. A value of 4.0 is given to fluorine, the most electronegative atom. The least electronegative element, francium, has an electronegativity value of 0.7. The values for all the other elements lie between these two extremes. Note that electronegativity values are pure numbers with no units.

■ Trends in the properties of the elements in group 1 and group 17

Trends in atomic and ionic radiiAt the right of the periodic table, the atomic radius is defined as half the distance between the nuclei of two covalently bonded atoms (Figure 3.19). For example, the bond length in a chlorine molecule (the distance between two chlorine nuclei) is 0.199 nm. Therefore the atomic radius of chorine is ½ × 199 = 99 pm (1 picometre [pm] = 10−12 m; 1 nanometre [nm] = 10−9 m). At the left of the periodic table, the atomic radius is that of the atom in the metal lattice (the metallic radius). For the noble gases the atomic radius is that of an isolated atom (the van der Waals’ radius).

In general the atomic radius of an atom is determined by the balance between two opposing factors:■ the shielding effect by the electrons of the inner shell(s) – this makes the atomic radius

larger. The shielding effect is the result of repulsion between the electrons in the inner shell and those in the outer or valence shell.

■ the nuclear charge (due to the protons) – this is an attractive force that pulls all the electrons closer to the nucleus. With an increase in nuclear charge, the atomic radius becomes smaller.

However, when moving down a group in the periodic table, there is an increase in the atomic radius as the nuclear charge increases (Tables 3.4 and 3.5 and Figures 3.20 and 3.21). This is the result of two factors:■ the increase in the number of complete electron shells between the outer (valence) electrons

and the nucleus■ the increase in the shielding effect of the outer electrons by the inner electrons.

metallic radius

a

b

c

covalent radius

van der Waals’ radius(for group 18)

r

r r

r

r

■ Figure 3.19 Atomic radius

4 Use graphs to plot the melting points, electron affinity values and first ionization energies for the halogens fluorine, chlorine, bromine and iodine (group 17 elements). Extrapolate the smooth curves to estimate the values for astatine. Compare your values with the values on the pages 7 and 8 of the IB Chemistry data booklet.


15


1

PERIODIC TRENDSPeriodic Trends in ionic radius The radii of cations and anions vary from the parent atoms from which they are formed in the following way.

The radii of cations are smaller than those of their parent atoms. The reason for this is that there are more protons than electrons in the cation so the valence electrons are more strongly attracted to the nucleus.

If an atom loses all its outer electrons, the radius of the resulting ion is much smaller than the atomic radius. This is because: • there is one fewer shell of electrons • there are fewer electrons in the positive ion than in the atom, so the electron–

electron repulsion is less, causing a further reduction in the radius.

31Electron affi nity (EA)

Table 2.6 Atomic and ionic radii

Group 1 Group 2 Group 3 Group 4 Group 5 Group 6 Group 7

Atomic Na Mg Al C N O F

Radius/pm 186 160 143 70 65 60 50

Ionic Na+ Mg2+ Al3+ N3− O2− F−

Radius/pm 95 65 50 171 140 136

The nuclear charge increases from 11 to 12 to 13. The force of attraction between the nucleus of the aluminium ion (13 protons) and its ten electrons is, therefore, greater than that between the nucleus of the other ions and their ten electrons. This causes the Al3+ ion to have the smallest radius.

Size of negative ionsA negative ion is always larger than its neutral atom (Figure 2.20). To form a negative ion, an atom must gain one or more electrons. The atom and the ion have the same atomic number, so forces of attraction remain the same. However, there is extra repulsion due to the increased number of electrons in the same shell. This causes the ion to expand, moving the electrons further from the nucleus until, once again, there is a balance between the forces of attraction and the forces of repulsion.

Na Na+ Cl Cl–

Electron af" nity (EA)Electron a! nity can be represented by the equation:

A(g) + e− → A−(g)

The negatively charged electron being added is brought towards the positively charged nucleus. There is a force of attraction between the two and, therefore, energy is released when the two are brought closer together.

The ! rst electron a" nity of oxygen is −142 kJ mol−1. This is the energy change, per mole, for the process:

O(g) + e− → O−(g)

The second electron a" nity can be represented by the equation:

A− (g) + e− → A2−(g)

The second electron a! nity of an element is always positive (endothermic) because energy is required to add an electron to an already negative ion. The incoming electron is repelled by the negative ion. Therefore, energy has to be supplied to bring the ion and the electron together.

Ions with the same electron structure are called isoelectronic.

Key termThe ! rst electron af! nity is the energy change when one electron is added to each atom in a mole of neutral gaseous atoms.

The values of " rst electron af" nities are negative (exothermic).

Figure 2.20 Relative sizes of atoms and ions

Key termThe second electron af! nity is the energy change when one electron is added to each ion in a mole of singly charged gaseous negative ions.

The second (and third) electron af" nities are always positive numbers. They are endothermic reactions.


3 Periodicity98

Trends in first ionization energyOn moving down a group, the atomic radius increases as additional electron shells are added. This causes the shielding effect to increase. The further the outer or valence shell is from the nucleus, the smaller the attractive force exerted by the protons in the nucleus. Hence, the more easily an outer electron can be removed and the lower the ionization energy. So, within each group, the first ionization energies decrease down the group. This is shown in Table 3.8 and Figure 3.24.

■ Figure 3.23 The relative sizes of the atoms and ions of group 1 metals

Li Na KNa+Li+ K+

■ Table 3.6 The variation of ionic radii in group 1

Ion Atomic number Ionic radius/pm

Li+ 3 68

Na+ 11 98

K+ 19 133

Rb+ 37 148

Cs+ 55 167

Fr+ 87 No data ■ Table 3.7 The variation of ionic radii in group 17

Ion Atomic number Ionic radius/pm

F– 9 133

Cl– 17 181

Br– 35 196

I– 353 219

At– 85 No data

■ Table 3.8 The variation of first ionization energy in group 1

Atom Atomic numberFirst ionization energy/kJ mol–1

Li 3 519

Na 11 494

K 19 418

Rb 37 402

Cs 55 376

Effective nuclear chargeAn alternative way to account for differences in ionization energy is to use the concept of effective nuclear charge.

Additional Perspectives

Atomic number0 10

LiNa

K RbCs

20 30 40 50 600

400

500

300

100

200

Firs

t ion

izat

ion

ener

gy/k

J m

ol–1

■ Figure 3.24 Bar graph showing the variation of first ionization energy in group 1

Trends in electronegativityElectronegativity values generally decrease down a group. Clear decreasing trends in electronegativity can be found in group 1 (the alkali metals, Table 3.9) and group 17 (the halogens, Table 3.10). Electronegativity can be interpreted as a measure of non-metallic or metallic character. Decreasing electronegativity down a group indicates a decrease in non-metallic character and an increase in metallic character.


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16


1

PERIODIC TRENDSFor ions with the same electron configuration (e.g. Na+, Mg2+ and Al3+), the ion with the greatest charge will have the smallest radius.


The radii of anions are larger than those of their parent atoms. A negative ion is always larger than its neutral atom. To form a negative ion, an atom must gain one or more electrons. The atom and the ion have the same atomic number, so forces of attraction remain the same. However, there is extra repulsion due to the increased number of electrons in the same shell. This causes the ion to expand, moving the electrons further from the nucleus.





Radius/pm 186 160 143 70 65 60 50


Radius/pm 95 65 50 171 140 136



Na Na+ Cl Cl–


A(g) + e− → A−(g)



O(g) + e− → O−(g)


A− (g) + e− → A2−(g)









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Radius/pm 95 65 50 171 140 136



Na Na+ Cl Cl–


A(g) + e− → A−(g)



O(g) + e− → O−(g)


A− (g) + e− → A2−(g)









17


1

PERIODIC TRENDS

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CHEMISTRY: FOR USE WITH THE IB DIPLOMA PROGRAMME STANDARD LEVEL

CH

APT

ER 3

PE

RIO

DIC

ITY

83

540

669

1310

519 900

494 736

418 590

402 548

632

636

376

661

669

531

648

653

760

653

694

770

716

699

762

762

724

841

757

745

887

736

803

866

745

732

891

908

866

1010

799

577

577

556

590

1090

786

762

707

716

1400

1060

966

833

703

1310(O– +844)

1000(S– +532)

941

870

812

1680

1260

1140

1010

920

2370

2080

1520

1350

1170

1040502

181

2.1

First ionizationenergy (kJ mol–1)

Electronegativity

1.0

0.9

0.8

0.8

0.7

0.7

1.5

1.2

1.0

1.0

0.9

0.9

1.3

1.2

1.1

1.1

1.5

1.4

1.3

1.6

1.6

1.5

1.6

1.8

1.7

1.5

1.9

1.9

1.8

2.2

2.2

1.8

2.2

2.2

1.8

2.2

2.2

1.9

1.9

2.4

1.6

1.7

1.9

2.0

1.5

1.6

1.7

1.8

2.5

1.8

1.8

1.8

1.8

3.0

2.1

2.0

1.9

1.9

3.5

2.5

2.4

2.1

2.0

4.0

3.0

2.8

2.5

2.2

510

La

Ac

H

Element

Li Be

Na Mg

K Ca

Rb Sr

Sc

Y

Hf

Ti

Zr

Ta

V

Nb

W

Cr

Mo

Re

Mn

Tc

Os

Fe

Ru

Ir

Co

Rh

Pt

Ni

Pd

Au

Cu

Ag

Hg

Zn

Cd

Tl

Ga

In

Pb

Ge

Sn

Bi

As

Sb

Po

Se

Te

At

Br

I

Rn

Kr

Al Si P S Cl Ar

B C N O F Ne

Xe

Cs Ba

Fr Ra

Increasing electronegativity and first ionization energy

Incr

easi

ng e

lect

rone

gativ

ity a

nd f

irst

ioni

zatio

n en

ergy

He

Figure 3.2.1 Electronegativity and first ionization energy values, showing trends within the periodic table. A similar table can be found in the IB Data Booklet © IBO 2007.

188

200

30

152 112

186 160

231 197

244 215

160

180

262

146

157

157

131

141

143

125

136

137

129

135

137

126

133

134

125

134

135

124

138

138

128

144

144

133

149

152

88

143

141

166

171

77

117

122

162

175

70

110

121

141

170

66

104

117

137

140

58

99

114

133

140217

270

154 (1–)

Atomicradius(10–12 m)

Ionicradius(10–12 m)

68 (1+)

98 (1+)

133 (1+)

148 (1+)

167 (1+)

30 (2+)

65 (2+)

94 (2+)

110 (2+)

34 (2+)

81 (3+)

93 (3+)

115 (3+)

90 (2+)68 (4+)

80 (4+)

81 (4+)

88 (2+)59 (5+)

70 (5+)

73 (5+)

63 (3+)

68 (4+)

68 (4+)

80 (2+)60 (4+)

76 (2+)64 (3+)

65 (4+)

67 (4+)

74 (2+)63 (3+)

86 (2+)

66 (4+)

72 (2+) 96 (1+)69 (2+)

126 (1+)

137 (1+)85 (3+)

74 (2+)

97 (2+)

127 (1+)110 (2+)

62 (3+)

81 (3+)

95 (3+)

53 (4+)272 (4–)

112 (2+)71 (4+)

120 (2+)84 (4+)

222 (3–)

245 (3–)

120 (3+)

202 (2–)

222 (2–)

196 (1–)

45 (3+) 42 (4+)271 (4–)

212 (3–) 190 (2–) 181 (1–)

16 (3+) 260 (4–) 171 (3–) 146 (2–) 133 (1–)

219 (1–)

220

La

Ac

H

Element

Li Be

Na Mg

K Ca

Rb Sr

Sc

Y

Hf

Ti

Zr

Ta

V

Nb

W

Cr

Mo

Re

Mn

Tc

Os

Fe

Ru

Ir Pt

Co

Rh

Ni

Pd

Au

Cu

Ag

Hg

Zn

Cd

Tl

Ga

In

Pb

Ge

Sn

Bi

As

Sb

Po

Se

Te

At

Br

I

Rn

Kr

Al Si P S Cl Ar

B C N O F Ne

Xe

Cs Ba

Fr Ra

Decreasing atomic and ionic radii

Incr

easi

ng a

tom

ic a

nd io

nic

radi

i

He

Figure 3.2.2 Atomic and ionic radii values showing trends within the periodic table. A similar table can be found in the IB Data Booklet © IBO 2007.

18


1

PERIODIC TRENDSSkill Check 1

CHEMISTRY: FOR USE WITH THE IB DIPLOMA PROGRAMME STANDARD LEVEL

CH

APT

ER 3

PE

RIO

DIC

ITY

87

7 For each of the following pairs, determine whether the atomic/ionic radius of the fi rst particle listed is larger than (>), the same size (=), or smaller than (<) the radius of the second particle.

Just as the physical properties exhibit a periodicity within groups and periods, so do the chemical properties of the elements and some of their compounds. The chemical properties relate to the electron arrangement of the elements.

The alkali metals are well known for their reactive nature. Their tendency to react sometimes violently with water means that they must be stored under oil. Potassium, in particular, is seldom found in secondary school laboratories because of its diffi culty in storage and its violent reaction with water.

The reactivity of the alkali metals with water increases down group 1. While lithium metal fl oats on the surface of the water and reacts slowly, producing some hydrogen gas, sodium reacts more violently, whizzing around on the surface of the water on a layer of hydrogen gas in what is sometimes described as ‘hovercraft’ motion.

2Li(s) + H2O(l) → Li2O(aq) + H2(g)

2Na(s) + H2O(l) → Na2O(aq) + H2(g)

The sodium can be forced to ignite by slowing its progress on the surface of the water by placing it on paper towel (fi gure 3.3.1). The white smoke in fi gure 3.3.1 is sodium oxide being carried off as a smoke. Potassium burns spontaneously in water, producing a violet fl ame.

2K(s) + H2O(l) → K2O(aq) + H2(g)

All three alkali metals produce an alkaline solution when they react with water.

Li2O(s) + H2O(l) → 2LiOH(aq)

Na2O(s) + H2O(l) → 2NaOH(aq)

K2O(s) + H2O(l) → 2KOH(aq)

Evidence for this alkaline solution can be seen in fi gure 3.3.1. Phenolphthalein indicator has been added to the water before the reaction with sodium and has turned pink due to the presence of sodium hydroxide.

The increase in reactivity can be explained by the decrease in electrostatic attraction between the outer-shell electron and the positive nucleus of the alkali metal. The further the outer shell is from the nucleus, the more easily it is lost in the reaction with water, and the more spectacular the result.

The increasing reactivity of the alkali metals can also be seen in their reaction with halogens. As lithium, sodium and potassium all lose electrons easily (are good reducing agents) and the halogens gain electrons easily (are good oxidizing agents, see chapter 10), these reactions are predictably violent.

First particle >, =, < Second particlea Chlorine atom (Cl) Chloride ion (Cl−)

b Aluminium ion (Al3+) Aluminium atom (Al)

c Calcium atom (Ca) Sulfur atom (S)

d Sodium ion (Na+) Fluoride ion (F−)

e Magnesium ion (Mg2+) Calcium ion (Ca2+)

f Sulfide ion (S2−) Potassium ion (K+)

3.3 CHEMICAL PROPERTIES OF ELEMENTS AND THEIR OXIDES

Trends in chemical properties within a group

3.3.1Discuss the similarities and differences in the chemical properties of elements in the same group. © IBO 2007

Figure 3.3.1 Sodium will burn with a yellow flame if its motion on the surface of water is stilled.

DEMO 3.2Reactions of group 1 and group 2 elements with water

19


1

PERIODIC TRENDSPeriodic trends in ionisation energy

The first ionisation energies of the elements show periodicity. The pattern from lithium to neon is repeated exactly with the elements sodium to argon. Apart from the insertion of the d-block elements, this pattern is seen again from potassium to rubidium.

17Ionisation energy

For the element calcium, the second ionisation energy is the energy change per mole for:

Ca+(g) → Ca2+(g) + e−

Test yourself5 Write an equation to show the ! fth ionisation energy of " uorine.

First ionisation energies of hydrogen to rubidiumThe # rst ionisation energies of the elements show periodicity. The pattern from lithium to neon is repeated exactly with the elements sodium to argon. Apart from the insertion of the d-block elements, this pattern is seen again from potassium to rubidium.

The trend down a group is for the # rst ionisation energy to decrease. Although the number of protons increases, so does the number of shielding electrons. The main factor is the increase in the atomic radius, making the outermost electron further from the nucleus and so easier to remove.

The general trend across a period is for the # rst ionisation energy to increase. However, there are a number of slight variations from this trend.

Figure 2.6 shows that there are maxima at each noble gas and minima at each group 1 metal. There are dips after the second and # fth elements in both periods 2 and 3.

TipThe second ionisation energy is always a positive number (an endothermic process).

0

200

400

600

800

1000

1200

1400

1600

H

He

N

CBe

B

Li Na

Ar

K

Mg

Kr

Rb

O

F

Ne

1800

2000

2200

2400

1st i

onis

atio

n en

ergy

/kJ

mol

–1

Atomic number50 10 15 20 25 30 35 40

Al

S

PZn

Ga

Se

As

Cl

Si Ge

Br

Figure 2.6 Variation of the ! rst ionisation energy with atomic number

The energy required for the process Ca(g) → Ca2+(g) + 2e− is the sum of the ! rst and second ionisation energies.


153

8Periodicity

Why the radii of atoms increase down a group Going down a group in the Periodic Table, the atoms of each element have one extra complete main shell of electrons compared with the one before. So, for example, in Group 1 the outer electron in sodium is in main shell 3, whereas in potassium it is in main shell 4. So going down the group, the outer electron main shell is further from the nucleus and the atomic radii increase.

First ionisation energyThe !rst ionisation energy is the energy required to convert a mole of isolated gaseous atoms into a mole of singly positively charged gaseous ions, that is, to remove one electron from each atom.

E(g) ➝ E+(g) + e– (g) where E stands for any element

The !rst ionisation energies also have periodic patterns. These are shown in Figure 4.

▲ Figure 4 The periodicity of !rst ionisation energies

!rst

ioni

satio

n en

ergy

/ kJ

mol

–1

0

2500

500

1500

2000

Li

Be

B

C

N

O

F

Ne

Mg

Al

Si

P

S

Cl

5 10 15

atomic number

0 20

1000H

NaK

Ar

Ca

He

Synoptic linkYou learnt about ionisation energy in Topic 1.6, Electron arrangements and ionistion energy.

+ The discovery of argonWhen the Periodic Table was !rst put forward, none of the noble gases had been discovered. The !rst, argon, was discovered by Scottish chemist William Ramsay. He noticed that the density of nitrogen prepared by a chemical reaction was 1.2505 g dm−3 while nitrogen prepared by removing oxygen and carbon dioxide from air had a density of 1.2572 g dm−3. He reasoned that there must be a denser impurity in the second sample, which he showed to be a previously unknown and very unreactive gas – argon. He later went on to discover the whole group of noble gases for which he won the Nobel (not noble!) Prize.

Chemists at the time had di#culty placing an unreactive gas of Ar, approximately 40 in the Periodic Table.

A suggestion was made that argon might be an allotrope of nitrogen, N3, analogous to the O3 allotrope of oxygen.

1 To how many signi!cant !gures were the two densities measured?

2 Using relative atomic masses from the Periodic Table, explain why argon is denser than nitrogen.

3 Suggest how oxygen could be removed from a sample of air.

4 What would be the Mr of N3 (to the nearest whole number)?

5 Suggest why chemists were reluctant to regard argon as an element.

HintRemember that electron levels are also referred to as electron shells.

20

Marginalizer Bilal Hameed


1

PERIODIC TRENDSThe general trend across a period is for the first ionisation energy to increase.

• As the effective nuclear charge increases from left to right across a period (nuclear charge increases and the shielding effectively remains the same), the valence electrons are pulled closer to the nucleus, so the attraction between the electrons and the nucleus increases. This makes it more difficult to remove an electron from the atom.

• Atomic radii decrease across a period – because the distance between the valence electrons and the nucleus decreases, it becomes more difficult to remove an electron from the atom.

However, there are a number of slight variations from this trend.

The trend down a group is for the first ionisation energy to decrease. Although the number of protons increases, so does the number of shielding electrons, hence keeping the effective nuclear charge constant. The main factor is the increase in the atomic radius, making the outermost electron further from the nucleus and so easier to remove.

154

Why the first ionisation energy increases across a periodAs you go across a period from left to right, the number of protons in the nucleus increases but the electrons enter the same main shell, see Figure 5. The increased charge on the nucleus means that it gets increasingly dif!cult to remove an electron.

Why the first ionisation energy decreases going down a groupThe number of !lled inner shells increases down the group. This results in an increase in shielding. Also, the electron to be removed is at an increasing distance from the nucleus and is therefore held less strongly. Thus the outer electrons get easier to remove going down a group because they are further away from the nucleus.

Why there is a drop in ionisation energy from one period to the nextMoving from neon in Period 0 (far right) with electron arrangement 2,8 to sodium, 2,8,1 (Period 1, far left) there is a sharp drop in the !rst ionisation energy. This is because at sodium a new main shell starts and so there is an increase in atomic radius, the outer electron is further from the nucleus, less strongly attracted and easier to remove.

[Ne]3s1

11+

[Ne]3s2

12+

[Ne]3s2 3p1

13+

[Ne]3s2 3p2

14+

[Ne]3s2 3p3

15+

[Ne]3s2 3p4

16+

[Ne]3s2 3p5

17+

[Ne]3s2 3p6

18+

Outer electrons are harder to remove as nuclear charge increases

▲ Figure 5 The electronic structures of the elements sodium to argon

The !rst ionisation energy generally increases across a period (see Figure 4), alkali metals like sodium, Na, and lithium, Li, have the lowest values and the noble gases (helium, He, neon, Ne, and argon, Ar) have the highest values.

The !rst ionisation energy decreases going down any group. The trends for Group 1 and Group 0 are shown dotted in red and green, respectively on the graph.

You can explain these patterns by looking at electronic arrangements (Figure 5).

Summary questions1 What happens to the size of atoms as you go from left to right across a

period? Choose from increase, decrease, no change.2 What happens to the !rst ionisation energy as you go from left to right

across a period? Choose from increase, decrease, no change.3 What happens to the nuclear charge of the atoms as you go left to right

across a period?4 Why do the noble gases have the highest !rst ionisation energy of all

the elements in their period?

HintFilled inner electron shells are said to shield electrons in the outer shell from the nuclear charge.

8.3 More trends in the properties of the elements of Period 3

3.2 Periodic trends 101

Ions of sodium, magnesium and aluminium are isoelectronic species (Table 3.15). The nuclear charge increases from the sodium ion to the aluminium ion. The higher nuclear charge pulls all the electron shells closer to the nucleus. Hence, the ionic radii decrease.

Similarly, the nuclear charge increases from the phosphide ion to the chloride ion. The higher nuclear charge causes the electron shells to be pulled closer to the nucleus. Again, the ionic radii decrease (Table 3.16).

Species Na+ Mg2+ Al3+

Nuclear charge +11 +12 +13

Number of electrons 10 10 10

Ionic radius/pm 98 65 45

The large increase in size from the aluminium ion to the phosphide ion is due to the presence of an additional electron shell. This causes a large increase in the shielding effect and as a result the ionic radius increases.

Trends in first ionization energyThe first ionization energies of the elements in period 3 are listed in Table 3.17. The general trend is an increase in first ionization energy across the periodic table. When moving across a period from left to right the nuclear charge increases but the shielding effect only increases slightly (since electrons enter the same shell). Consequently, the electron shells are pulled progressively closer to the nucleus and as a result first ionization energies increase.

Element Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine

First ionization energy/kJ mol–1

494 736 577 786 1060 1000 1260

However, the increase in first ionization energy is not uniform and there are two decreases – between magnesium and aluminium and between phosphorus and sulfur. These decreases can only be explained by reference to sub-shells and orbitals.

The first ionization energy of aluminium is lower than that of magnesium, even though aluminium has a smaller atomic radius. The decrease in first ionization energy from magnesium (1s2 2s2 2p6 3s2) to aluminium (1s2 2s2 2p6 3s2 3p1) occurs because the electrons in the filled 3s orbital are more effective at shielding the electron in the 3p orbital than they are at shielding each other. Therefore less energy is needed to remove a single 3p electron than to remove a paired 3s electron.

The first ionization energy of sulfur (1s2 2s2 2p6 3s2 3p2 3p1 3p1) is less than that of phosphorus (1s2 2s2 2p6 3s2 3p13p1 3p1) because less energy is required to remove an electron from the 3p4 orbitals of sulfur than from the half-filled 3p orbitals of phosphorus. The presence of a spin pair of electrons results in greater electron repulsion compared to two unpaired electrons in separate orbitals.

■ Trends in electronegativity valuesThe electronegativities of the elements in period 3 are listed in Table 3.18. The general trend is an increase in first ionization energy across the periodic table. When moving across a period from left to right the nuclear charge increases but the shielding effect only increases slightly (since electrons enter the same shell). Consequently, the electron shells are pulled progressively closer to the nucleus and as a result electronegativity values increase.

Element Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine

Electronegativity 0.9 1.3 1.6 1.9 2.2 2.6 3.2

Generally, the electronegativity values of chemical elements increase across a period and decrease down a group (Figure 3.28). This observation can be used to compare the relative electronegativity values of two elements in the periodic table. To do this, find the positions of

Species P3– S2– Cl–

Nuclear charge +15 +16 +17

Number of electrons 18 18 18

Ionic radius/pm 212 190 181

■ Table 3.16 Atomic data for phosphide, sulfide and chloride ions

■ Table 3.15 Atomic data for sodium, magnesium and aluminium ions

■ Table 3.17 First ionization energies for the elements in period 3

■ Table 3.18 Electronegativity values for the elements in period 3


121

Explaining periodic patterns

Down a groupDown a group, the ionisation energy decreases.

The first ionisation energies (kJ mol–1) of the elements in Group 1 of the periodic table are shown in Table 8.6.

Table 8.6 First ionisation energies of the Group 1 elements.

Element Li Na K Rb Cs

First ionisation energy/kJ mol–1 520 496 419 403 376

Factor 1 Atomic radius – increases down the group and the outer electron is further from the nucleus. This tends to decrease the ionisation energy as we go down the group.

Factor 2 Nuclear charge – increases down the group. This tends to increase the ionisation energy as we go down the group.

Factor 3 Shielding effect – increases down the group. This reduces the effect of the nucleus. This tends to decrease the ionisation energy as we go down the group.

Factors 1 and 3 outweigh factor 2 as ionisation energy decreases down a group.

Across a periodAcross a period, ionisation energy increases. This is illustrated in Table 8.7.

Table 8.7 First ionisation energies of elements in Period 2.

Element Li Be B C N O F Ne

First ionisation energy / kJ mol–1

520 900 801 1086 1402 1314 1681 2081

Electron structure 1s22s1 1s22s2

full sub shell1s22s22p1 1s22s22p2 1s22s22p3

½ full sub-shell1s22s22p4 1s22s22p5 1s22s22p6

full shell

We can also explain this in terms of the three factors.

Factor 1 Atomic radius – decreases across the period so the outer electrons are closer to the nucleus. This tends to increase the ionisation energy as we go across the period.

Factor 2 Nuclear charge – increases across the period. This tends to increase the ionisation energy as we go across the period.

Factor 3 Shielding effect – generally remains the same across a period and therefore has little effect on the nucleus.

Factors 1 and 2 both indicate that ionisation energy will increase across the period.

There is a general increase across a period but there are slight decreases after the second and the fifth element.

! Elements in Group 13 have an electron structure of s2p1 and the s electrons provide a slightly greater shielding of the p electron, which is therefore lost a little more readily.

! Elements in Group 15 have a half-filled set of p orbitals but in Group 16 there is a pair of p electrons and the repulsion between this pair of electrons is sufficient to make the elements in Group 16 ionise slightly more readily.


28

Both these cases, which go against the expected trend, are evidence that con!rms the existence of s- and p-sub-shells. These were predicted by quantum theory and the Schrödinger equation.

Trends in ionisation energies down a group in the Periodic TableFigure 5 shows that there is a general decrease in !rst ionisation energy going down Group 2, and the same pattern is seen in other groups. This is because the outer electron is in a main shell that gets further from the nucleus in each case.

In Figure 2, notice the small drop between phosphorus (1s2, 2s2, 2p6, 3s2, 3p3) and sulfur (1s2, 2s2, 2p6, 3s2, 3p4). In phosphorus, each of the three 3p orbitals contains just one electron, while in sulfur, one of the 3p orbitals must contain two electrons. The repulsion between these paired electrons makes it easier to remove one of them, despite the increase in nuclear charge, see Figure 4.

Summary questions1 State why the second

ionisation energy of any atom is larger than the !rst ionisation energy.

2 Sketch a graph similar to Figure 1 of the successive ionisation energies of aluminium (electron arrangement 2,8,3).

3 An element X has the following values (in kJ mol–1) for successive ionisation energies: 1093, 2359, 4627, 6229, 37 838, 47 285. a Identify which group in the

Periodic Table it is in.b Explain your answer to a.

Going down a group, the nuclear charge increases. At !rst sight you might expect that this would make it more dif!cult to remove an electron. However, the actual positive charge ‘felt’ by an electron in the outer shell is less than the full nuclear charge. This is because of the effect of the inner electrons shielding the nuclear charge.

▲ Figure 5 The first ionisation energies of the elements of Group 2

!rst

ioni

satio

n en

ergy

/ kJ

mol

–1

900

850

800

750

700

650

600

550

500

Mg

Ca

Sr

Ba

atomic number0 10 20 30 40 50 60

Be

▲ Figure 4 Electron arrangements of phosphorus and sulfur

orbitals (sub-shells) in phosphorus

1s 2s2p

3s3p

orbitals (sub-shells) in sulfur

easierto lose

1s 2s2p

3s3p

1.6 Electron arrangements and ionisation energy

21


1

PERIODIC TRENDSWhy there is a drop in ionisation energy from one period to the nextMoving from neon in Period 2 Group 18 (far right) with electron arrangement 2,8 to sodium, 2,8,1 (Period 3 Group 1 far left) there is a sharp drop in the first ionisation energy. This is because at sodium a new main shell starts and so there is an increase in atomic radius, the outer electron is further from the nucleus, less strongly attracted and easier to remove.

A closer look at ionisation energies of Period 3

The graph of first ionisation energy against atomic number across a period is not smooth. It shows that:

the first ionisation energy actually drops between Group 2 and Group 13, so that aluminium has a lower ionisation energy than magnesium

the ionisation energy drops again slightly between Group 15 (phosphorus) and Group 16 (sulfur).

Similar patterns occur in other periods. You can explain this if you look at the electron arrangements of these elements.

155

Learning objectives:➔ Explain why the increase in

ionisation energies across a period is not regular.

➔ Describe how successive ionisation energies explain electron arrangements.


This chapter revisits the trends in ionisation energies !rst dealt with in Topic 1.6, in the context of periodicity. The graph of !rst ionisation energy against atomic number across a period is not smooth. Figure 1 below shows the plot for Period 3.

HintIonisation energies are sometimes called ionisation enthalpies.

atomic number

!rst

ioni

satio

n en

ergy

/ kJ

mol

–1

100

250

500

750

1000

1250

1500

1750

11 12 13 14 15

Group 1Na

Group 2Mg

Group 3Al

Group 4Si

Group 5P

Group 6S

Group 7Cl

Group 0Ar

16 17 18

▲ Figure 1 Graph of !rst ionisation energy against atomic number for the elements of Period 3

It shows that:

• the !rst ionisation energy actually drops between Group 2 and Group 3, so that aluminium has a lower ionisation energy than magnesium

• the ionisation energy drops again slightly between Group 5 (phosphorus) and Group 6 (sulfur).


The drop in first ionisation energy between Groups 2 and 3For the !rst ionisation energy:

• magnesium, 1s2 2s2 2p6 3s2, loses a 3s electron

• aluminium, 1s2 2s2 2p6 3s2 3p1, loses the 3p electron.

The p-electron is already in a higher energy level than the s-electron, so it takes less energy to remove it, see Figure 2.

complete removal

magnesium 1s2 2s2 2p6 3s2

ener

gy

1s

2s

3s

2p

3p

1st IE

complete removal

aluminium 1s2 2s2 2p6 3s2 3p1

ener

gy

1s

2s

3s

2p

3p

1st IE

▲ Figure 2 The !rst ionisation energies of magnesium and aluminium (not to scale)

8.4 A closer look at ionisation energies22


1

PERIODIC TRENDSThe drop in first ionisation energy between Groups 2 and 3 For the first ionisation energy:

• magnesium, 1s2 2s2 2p6 3s2 , loses a 3s electron

• aluminium, 1s2 2s2 2p6 3s2 3p1 , loses the 3p electron

The p-electron is already in a higher energy level than the s-electron, so it takes less energy to remove it

The drop in first ionisation energy between Groups 15 and 16 An electron in a pair will be easier to remove that one in an orbital on its own because it is already being repelled by the other electron.

• phosphorus, 1s2 2s2 2p6 3s2 3p3, as no paired electrons in a p-orbital because each p-electron is in a different orbital

• sulfur, 1s2 2s2 2p6 3s2 3p4, has two of its p-electrons paired in a p-orbital so one of these will be easier to remove than an unpaired one due to the repulsion of the other electron in the same orbital.

155







atomic number

!rst

ioni

satio

n en

ergy

/ kJ

mol

–1

100

250

500

750

1000

1250

1500

1750

11 12 13 14 15

Group 1Na

Group 2Mg

Group 3Al

Group 4Si

Group 5P

Group 6S

Group 7Cl

Group 0Ar

16 17 18


It shows that:








complete removal


ener

gy

1s

2s

3s

2p

3p

1st IE

complete removal


ener

gy

1s

2s

3s

2p

3p

1st IE


8.4 A closer look at ionisation energies

155







atomic number

!rst

ioni

satio

n en

ergy

/ kJ

mol

–1

100

250

500

750

1000

1250

1500

1750

11 12 13 14 15

Group 1Na

Group 2Mg

Group 3Al

Group 4Si

Group 5P

Group 6S

Group 7Cl

Group 0Ar

16 17 18


It shows that:








complete removal


ener

gy

1s

2s

3s

2p

3p

1st IE

complete removal


ener

gy

1s

2s

3s

2p

3p

1st IE



156


The drop in first ionisation energy between Groups 5 and 6An electron in a pair will be easier to remove that one in an orbital on its own because it is already being repelled by the other electron. As shown in Figure 3:

• phosphorus, 1s2 2s2 2p6 3s2 3p3, has no paired electrons in a p-orbital because each p-electron is in a different orbital

• sulfur, 1s2 2s2 2p6 3s2 3p4, has two of its p-electrons paired in a p-orbital so one of these will be easier to remove than an unpaired one due to the repulsion of the other electron in the same orbital.

orbitals (sub-shells) in phosphorus

1s 2s2p

3s3p

orbitals (sub-shells) in sulfur

easierto lose

1s 2s2p

3s3p

▲ Figure 3 Electron arrangements of phosphorus and sulfur

Successive ionisation energiesIf you remove the electrons from an atom one at a time, each one is harder to remove than the one before. Figure 4 is a graph of ionisation energy against number of electrons removed for sodium, electron arrangement 2,8,1.

You can see that there is a sharp increase in ionisation energy between the !rst and second electrons. This is followed by a gradual increase over the next eight electrons and then another jump before the !nal two electrons. Sodium, in Group 1 of the Periodic Table, has one electron in its outer main shell (the easiest one to remove), eight in the next main shell and two (very hard to remove) in the innermost main shell.

Figure 5 is a graph of successive ionisation energies against number of electrons removed for aluminium, electron arrangement 2,8,3.

It shows three electrons that are relatively easy to remove – those in the outer main shell – and then a similar pattern to that for sodium.

If you plotted a graph for chlorine, the !rst seven electrons would be relatively easier to remove than the next eight.

This means that the number of electrons that are relatively easy to remove tells us the group number in the Periodic Table. For example, the values of 906, 1763, 14 855, and 21 013 kJ mol–1 for the !rst !ve ionisation energies of an element, tell us that the element is in Group 2. This is because the big jump occurs after two electrons have been removed.

0

electronsin shell 1

electrons in shell 2

electronsin shell 3

total number of electrons removed

log

IE

1 2 3 4 5 6 7 8 9 10 11

▲ Figure 4 Graph of successive ionisation energies against number of electrons removed for sodium. Note that the log of the ionisation energies is plotted in order to "t the large range of values onto the scale

0

electronsin shell 1

electrons in shell 2

electronsin shell 3

total number of electrons removed

log IE

1 2 3 4 5 6 7 8 9 10 111213

▲ Figure 5 Graph of successive ionisation energies against number of electrons removed for aluminium

Summary questions1 Write the electron

arrangement in the form 1s2… for:a berylliumb boron.

2 If one electron is lost from for the following atoms, from what main shell does it come?a berylliumb boron

3 Why is the !rst ionisation energy of boron less than that of beryllium?

23


1

PERIODIC TRENDS

Skill Check 2 The 1st ionisation energies of several elements with consecutive atomic numbers are shown in the graph below. The letters are not the symbols of the elements. a. Which of the elements A to I belong to Group I in the Periodic Table? Explain your answer. b. Which of the elements A to I could have the electronic configuration 1s2 2s2 2p6 3s2? c. Explain the rise in 1st ionisation energy between element E and element G. d. Estimate the 1st ionisation energy of element J.


10 Explain why the hydrogen atom, H(g) has a smaller value of first ionization energy than the helium ion, He+(g).

A similar explanation also accounts for the first decrease observed in period 3 for the elements magnesium and aluminium. The decrease in first ionization energy from magnesium (1s22s22p63s2) to aluminium (1s22s22p63s23p1) arises largely because the electrons in the filled 3s orbital are more effective at shielding the electron in the 3p orbital than they are at shielding each other.

The second decrease in first ionization energy in period 2 occurs between nitrogen (1s22s22px

12py12pz

1) and oxygen (1s22s22px22py

12pz1). The three valence (outer) electrons in

the nitrogen atom are in three separate orbitals. This is in accordance with Hund’s rule, which states that every orbital in a sub-shell is singly occupied with one electron before any one orbital is doubly occupied. However, in the oxygen atom two electrons are in the same 2p orbital. The two electrons in the same orbital experience severe repulsion. This electron–electron repulsion makes it easier to remove one of these 2px electrons than an unpaired electron from a half-filled 2pz orbital. Hence, the decrease in first ionization energy between nitrogen and oxygen is due to the additional repulsion present in the 2p4 configuration of the oxygen atom (Figure 12.18).

A similar explanation accounts for the decrease in first ionization energy observed between phosphorus and sulfur in period 3. The first ionization energy of sulfur (1s22s22p63s23px

23py13pz

1) is less than that of phosphorus (1s22s22p63s23px

13py13pz

1) because less energy is required to remove an electron from the 3p4 orbitals of sulfur than from the half-filled 3p orbitals of phosphorus.

The patterns of first ionization energies across periods 2 and 3 are identical except that all the corresponding ionization energies for period 2 are higher. This is because the electrons being removed are in a second shell closer to the nucleus, compared to the third shell for the period 3 elements. The outer electrons in period 2 experience a higher nuclear charge than those in period 3.

Explaining the trends and discontinuities in first ionization energy across a periodFigure 12.19 shows the relationship between first ionization energy and atomic number. The general increase in first ionization energy across a period is because the nuclear charge (proton number) increases across the period; shielding by other electrons only increases slightly and hence there is a greater attraction (higher effective nuclear charge) for the electrons.

oxygen atom, O1s 2s 2p

oxygen ion, O+

1s 2s 2p

nitrogen atom, N1s 2s 2p

nitrogen ion, N+

1s 2s 2p

Q Figure 12.18 Orbital notation for nitrogen and oxygen atoms and their unipositive ions

Q Figure 12.19 The relationship between first ionization energy and atomic number

Ioni

zatio

n en

ergy

Atomic number

s sub-shell

filling up thep sub-shell filling up the

p sub-shell

half-filledp sub-shell

half-filledp sub-shell

s sub-shell s sub-shell

H

He

Li

Ne

Na

Ar

K

The ‘drops’ between groups 2 and 13 occur because an electron is removed from a p orbital which is higher in energy than an s orbital and hence is easier to remove. There is also an increase in shielding.

The ‘drops’ between groups 15 and 16 occurs because there are paired electrons in a p orbital, so less energy is required to remove one. This is due to the increased electron–electron repulsion.

There is a large drop at the end of each period because an electron shell has been completed, which results in a large increase in shielding. Hence there is less attraction for the electrons by the nucleus.

For the s- and p-blocks the increase in nuclear charge across a period has a large effect on the outer shell (valence) electrons because the inner shielding only increases slightly and


3 a What do you understand by the term atomic orbital? [1]b Draw diagrams to show the shape of:

i an s orbital [1]ii a p orbital. [1]

c Element X has the electronic confi guration 1s2 2s2 2p6 3s2 3p6 3d8 4s2.i Which block in the Periodic Table does element X belong to? [1]ii State the maximum number of electrons in a d sub-shell. [1]iii Element X forms an ion of type X2+. Write the full electronic confi guration for this ion using 1s2 notation. [1]iv Write the symbol for the sub-shell which begins to fi ll after the 3d and 4s are completely full. [1]

Total = 7

4 ! e 1st ionisation energies of several elements with consecutive atomic numbers are shown in the graph below. ! e letters are not the symbols of the elements.

2000

1600

1200

1800

400

0A B C D

Element

Firs

t ion

isatio

n en

ergy

/kJm

ol–1

E F G H I J

a Which of the elements A to I belong to Group I in the Periodic Table? Explain your answer. [3]b Which of the elements A to I could have the electronic confi guration 1s2 2s2 2p6 3s2? [1]c Explain the rise in 1st ionisation energy between element E and element G. [4]d Estimate the 1st ionisation energy of element J. [2]

46 3 Electrons in atoms

24


1

PERIODIC TRENDSSkill Check 3 The first ionisation of elements sodium to argon is shown below.

A Explain why the general trend from sodium to argon is upwards but why the value for sulfur is less than that for phosphorus. B Mark on the graph where the value for potassium would be. C Explain why the value for the second ionisation of sodium is very much larger than that of its first ionisation.

Skill Check 4 For each of the following pairs, state which element has the higher first ionisation energy and explain your answer: A Mg and Al

B Mg and Ca

C Ne and Na

2 Atomic structure and the periodic table (Topic 1)36

Exam practice questions 1 a) Define the term first ionisation

energy. (2)b) This part is about four sets of ionisation

energies:i) Which are the values of the successive

ionisation energies for an element in group 4 of the periodic table?A 496, 738, 578, 789, 1012, 1000B 578, 1817, 2745, 11 578, 14 831, 18 378C 1086, 2353, 4621, 6223, 37 832, 47 278D 1314, 1000, 941, 869, 812 (1)

ii) Which are the values for the first ionisation energies of consecutive elements in the same period?A 496, 738, 578, 789, 1012, 1000B 578, 1817, 2745, 11 578, 14 831,

18 378C 1086, 2353, 4621, 6223, 37 832,

47 278D 1314, 1000, 941, 869, 812 (1)

iii) Which are the values for the first ionisation energies of elements in the same group, as the group is descended?A 496, 738, 578, 789, 1012, 1000B 578, 1817, 2745, 11 578, 14 831,

18 378C 1086, 2353, 4621, 6223, 37 832,

47 278D 1314, 1000, 941, 869, 812 (1)

c) i) Define the term electron affinity. (2)ii) Which is the equation that relates to

the first electron affinity of chlorine?

A 12 Cl2(g) + e− → Cl−(g)

B Cl(g) + e− → Cl−(g)C Cl(g) − e− → Cl−(g)D Cl2(g) + 2e− → 2Cl−(g) (1)

(Total 8 marks)

2 This question is about mass spectrometry.a) A mass spectrometer can be used to find the

percentage composition of the isotopes of an element. Explain how the following are achieved in a mass spectrometer:i) ionisation (1)ii) acceleration (1)iii) deflection (1)

b) Analysis of a sample of iron in a mass spectrometer gave the following results:

Calculate the relative atomic mass of iron to two decimal places. (2)

c) The mass spectrum of bromine has lines at mz values of 158, 160 and 162, but none at 159 or 161. What causes the line at 160?A (80Br−80Br)+

B (80Br−80Br)−

C (79Br−81Br)−

D (79Br−81Br)+ (1)d) State and outline one modern use of mass

spectrometry. (3)(Total 9 marks)

3 The first ionisation of elements sodium to argon is shown below.

Firs

t io

nisa

tion

ene

rgy

Na Mg Al Si P S Cl Ar K

a) Explain why the general trend from sodium to argon is upwards but why the value for sulfur is less than that for phosphorus. (5)

b) Mark on the graph where the value for potassium would be. (1)

c) Explain why the value for the second ionisation of sodium is very much larger than that of its first ionisation. (2)

(Total 8 marks)

Isotope Relative isotopic mass %54Fe 53.94 5.9456Fe 55.93 91.7857Fe 56.94 2.28


25


05 periodic trends note - chemwithbilal.com

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