the periodic table 4 - pedagogics.ca · 2012-02-20 · 144 actinides lanthanides transition...

1434 THE PERIODIC TABLE

The periodic table 44.1 The periodic tableThe elements in the periodic table are arranged in order of atomic number, starting with hydrogen, which has atomic number 1. The groups are arranged in the vertical columns in the periodic table and the periods are arranged in the horizontal rows (Figure 4.1).

Most of the elements in the periodic table are metals – these are shown in yellow in Figure 4.1. The rest of the elements are non-metals, although some elements, such as Si, Ge and Sb, which are near the dividing line between metals and non-metals, are sometimes classi� ed as semi-metals or metalloids.

Learning objectives

• Understand how the elements in the periodic table are arranged

• Understand the terms group and period

• Understand how the electronic con� guration of an element relates to its position in the periodic table

26Fe

3group number

1

1

2

3

4

peri

od n

umbe

r

5

7

6

2 4 5 6 7

0

27Co

28Ni

29Cu

30Zn

31Ga

32Ge

33As

34Se

35Br

36Kr

52Te

51Sb

50Sn

49In

48Cd

47Ag

79Au

78Pt

77Ir

76Os

75Re

74W

73Ta

72Hf

57La

56Ba

55Cs

81TI

82Pb

83Bi

84Po

85At

80Hg

86Rn

46Pd

45Rh

44Ru

43Tc

42Mo

41Nb

40Zr

39Y

38Sr

37Rb

89Ac

88Ra

87Fr

53I

54Xe

13AI

14Si

15P

16S

18Ar

17CI

6C

5B

10Ne

2He

8O

7N

9F

25Mn

24Cr

23V

22Ti

1H

20Ca

19K

12Mg

11Na

4Be

3Li

21Sc

Figure 4.1 The periodic table. Hydrogen is sometimes shown, as here, in group 1. It may also be placed in the middle of the periodic table and not in any group.

The symbols of the elements that are solid at room temperature and pressure are shown in black in Figure 4.1, whereas those that are gases are in blue and liquids are in red.

Some of the groups in the periodic table are given names. Commonly used names are shown in Figure 4.2. The noble gases are sometimes also called the ‘inert gases’.

In the periodic table shown in Figure 4.1 it can be seen that the atomic numbers jump from 57 at La (lanthanum) to 72 at Hf (hafnium). This is because some elements have been omitted. These are the lanthanide elements. The actinide elements, which begin with Ac (actinium), have also been omitted from Figure 4.1. Figure 4.2 shows a long form of the periodic table, showing these elements as an integral part.

Hydrogen is the most abundant element in the universe: about 90% of the atoms in the universe are hydrogen. Major uses of hydrogen include making ammonia and hydrogenation of unsaturated vegetable oils, both discussed in later chapters.

144

ACTINIDES

LANTHANIDES

TRANSITION

ELEMENTS

HA

LOG

ENS

ALK

ALI

MET

ALS

ALK

ALI

NE

EART

H M

ETA

LS

NO

BLE

GA

SES

Figure 4.2 The long form of the periodic table, with the names of some of the groups. Hydrogen, though sometimes placed in group 1, does not count as an alkali metal.

The modern periodic table has developed from one originally conceived by Russian chemist Dmitri Mendeleev in 1869.

Mendeleev suggested that the elements were arranged in order of atomic weight (what we would now call relative atomic mass) and produced a table in which elements with similar chemical properties were arranged in horizontal groups. Mendeleev took several risks when presenting his data – he suggested that some elements had not been already discovered and left spaces for them in his table. Not only did he leave spaces but he also predicted the properties of these unknown elements – he made his hypotheses falsi� able, which added great weight to his theory. The predictions he made were later found to be extremely accurate – the mark of a good theory is that it should be able to be used to predict results that

can be experimentally con� rmed or refuted. He also suggested that the atomic weight of some elements were incorrect – he realised that Te belonged in the same group as O, S and Se, but its atomic weight was higher than iodine and so it should be placed after iodine. Instead of then abandoning his theory, he questioned the accuracy of the atomic weight of tellurium and placed it before iodine. This is, of course, the correct place for Te, but Mendeleev’s assumption that the atomic weight was lower than that of iodine was not correct.

German Chemist Julius Lothar Meyer was working on the arrangements of elements at the same time as Mendeleev and came to very similar conclusions – why is Mendeleev remembered as the father of the modern periodic table rather than Meyer?

Mendeleev puzzled over the arrangement of elements in the periodic table until he had a dream in which he claims to have seen the arrangement. Kekulé also came

up with the ring structure of benzene after a dream. Does it matter how a scientist comes up with a hypothesis? What is the di� erence between a scienti� c and a non-scienti� c hypothesis? Is it the origins of the hypothesis or the fact that it can be tested experimentally (is falsi� able) that makes it scienti� c?


The periodic table and electronic confi gurationsThe group number of an element indicates the number of electrons in its outer shell (the highest main energy level). Thus, all elements in group 1 have one electron in their outer shell, all elements in group 2 have two electrons in their outer shell, etc. The period number indicates the number of shells (main energy levels) in an atom.

26Fe

3numberof electronsin outer shell

1

1

2

3

4

num

ber

of s

hells

5

7

6

2 4 5 6 7

0

27Co

28Ni

29Cu

30Zn

31Ga

32Ge

33As

34Se

35Br

36Kr

52Te

51Sb

50Sn

49In

48Cd

47Ag

79Au

78Pt

77Ir

76Os

75Re

74W

73Ta

72Hf

57La

56Ba

55Cs

81TI

82Pb

83Bi

84Po

85At

80Hg

86Rn

46Pd

45Rh

44Ru

43Tc

42Mo

41Nb

40Zr

39Y

38Sr

37Rb

89Ac

88Ra

87Fr

53I

54Xe

13AI

14Si

15P

16S

18Ar

17CI

6C

5B

10Ne

2He

8O

7N

9F

25Mn

24Cr

23V

22Ti

1H

20Ca

19K

12Mg

11Na

4Be

3Li

21Sc

Learning objectives

• De� ne � rst ionisation energy and electronegativity of an element

• Understand trends in atomic radius, � rst ionisation energy and electronegativity across a period

• Understand trends in atomic radius, � rst ionisation energy and electronegativity down group 1 and group 7

• Explain the variation of melting points for elements across period 3 and down group 1 and group 7

Application 1Application 1

Let us consider P: this is in period 3 and group 5 and thus has three shells of electrons and � ve electrons in the outer shell. Selenium (Se) is in period 4 and group 6; it therefore has four shells of electrons and six electrons in the outer shell. The noble gases (group 0) have either two (He) or eight electrons in their outer shell.

4.2 Physical properties

Variation of properties down a group and across a periodIn the next few sections we will consider how various physical properties vary down a group or across a period in the periodic table.

ElectronegativityIn a covalent bond between two di� erent atoms, the atoms do not attract the electron pair in the bond equally. How strongly the electrons are attracted depends on the size of the individual atoms and their nuclear charge.

Electronegativity decreases down a group – this is because the size of the atom increases down a group. Consider hydrogen bonded to either F or Cl (Figure 4.3). The bonding pair of electrons is closer to the F nucleus in HF than it is to the Cl nucleus in HCl. Therefore the electron pair is more strongly attracted to the F nucleus in HF and F has a higher electronegativity than Cl.

Four elements are named after the small village of Ytterby in Sweden: yttrium, terbium, erbium and ytterbium.

146

Chlorine's higher nuclear charge does not make it more electronegative than � ourine because the shielding from inner shells (shown with blue shading in Figure 4.3) increases from F to Cl such that the e� ective nuclear charge felt by the bonding electrons is approximately the same in each case (+7 in each case, if shielding were perfect).

CI17+

F9+

H1+

H1+

a

b

shielding

attraction forthese electrons

greater distancebetween nucleusand bonding pairof electrons

CI17+

F9+

H1+

H1+

a

b

shielding

attraction forthese electrons

greater distancebetween nucleusand bonding pairof electrons

Electronegativity is a measure of the attraction of an atom in a molecule for the electron pair in the covalent bond of which it is a part.

Figure 4.3 Hydrogen bonded to (a) fl uorine and (b) chlorine.

N7+

H1+

a

b

shielding

F9+

H1+

shielding

attractionfor theseelectrons

highernuclearcharge


N7+

H1+

a

b

shielding

F9+

H1+

shielding


highernuclearcharge


Figure 4.4 Hydrogen bonded to (a) nitrogen and (b) fl uorine.

The full de� nition of � rst ionisation energy is: the energy required to remove one electron from each atom in one mole of gaseous atoms under standard conditions.

Electronegativity increases across a period – the reason for this is the increase in nuclear charge across the period with no signi� cant change in shielding. The shielding remains approximately constant because atoms in the same period have the same number of inner shells.

Thus, if an N–H bond is compared with an F–H bond (Figure 4.4), the electrons in the N–H bond are attracted by the seven protons in the nucleus, but the electrons in the F–H bond are attracted by the nine protons in the F nucleus. In both cases the shielding is approximately the same (because of two inner shell electrons).

First ionisation energyThe � rst ionisation energy for an element is the energy required to remove the outermost electron from a gaseous atom: that is, the energy for the following process:

M(g) → M+(g) + e−

Variation in fi rst ionisation energy down a groupThe decrease in � rst ionisation energy down a group can be seen in Figure 4.5. This is because the size of the atom increases so that the outer electron is further from the nucleus and therefore less strongly attracted by the nucleus (Figure 4.6).


Down any group in the periodic table the � rst ionisation energy decreases.

Although the nuclear charge also increases down a group, this is largely balanced out by an increase in shielding down the group, as there are more electron energy levels (shells). It is the increase in size that governs the change in � rst ionisation energy.

Variation in fi rst ionisation energy across a period

The general trend is that � rst ionisation energy increases from left to right across a period. This is because of an increase in nuclear charge across the period.

The nuclear charge increases from Na (11+) to Ar (18+) as protons are added to the nucleus. The electrons are all removed from the same main energy level (third shell) and electrons in the same energy level do not shield each other very well. Therefore the force on the outer electrons due to the nucleus increases from left to right across the period and the outer

Firs

t io

nisa

tion

ener

gy /

kJ m

ol –1

Li

Na

RbK

Cs

Group 1

1600

1400

1200

1000

800

600

4000

Figure 4.5 First ionisation energy for group 1. K19+

electron closerto the nucleus

Li3+

Figure 4.6 Potassium has a lower fi rst ionisation energy than lithium.

1600

1400

1200

1000

800

600

400

Firs

t io

nisa

tion

ener

gy /

kJ m

ol –1

0

Na

Mg

AI

Si

P S

CI

Ar

Period 3

Figure 4.7 The variation in fi rst ionisation energy across period 3 in the periodic table.

The increase in � rst ionisation energy (Figure 4.7) can also be explained in terms of the e� ective nuclear charge felt by the outer electron in argon being higher. The e� ective nuclear charge felt by the outer electron in a sodium atom would be 11 (nuclear charge) − 10 (number of inner shell electrons), i.e. 1+ if shielding were perfect. The e� ective nuclear charge felt by the outer electrons in an argon atom would be 18 (nuclear charge) − 10 (number of inner shell electrons), i.e. 8+ if shielding were perfect.

There are two exceptions to the general increase in � rst ionisation energy across a period, and these are discussed on pages 77 and 78. This is required only by Higher Level students.

148

Atomic radiusThe 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.

11+ 18+

Na Ar

atomic radius

Figure 4.8 The atomic radius of chlorine atoms in a molecule. K

Li

K

LiAtomic 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.

Extension

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


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.

Ionic radiusThe 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−.

80

100

120

140

160

180

200A

tom

ic r

adiu

s / p

m

AI

SiP

S CI

Na

Mg

Period 3

Figure 4.9 The variation in atomic radius across period 3. No atomic radius is shown for argon, as it does not form covalent bonds and the internuclear distance between atoms bonded together cannot be measured.

Extension

Although it is not possible to measure an atomic radius for Ar, it is possible to measure a value for the van der Waals’ radius of this element.

11+

Na

17+

CI

1 pm = 1 × 10−12 m

150

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.

0

250

200

300

150

100

50

Radi

us /

pm

Li

Na

CsRb

K

Li+

Na+

Cs+

Rb+

K+

Group 1

Figure 4.10 Atomic and ionic radii for the alkali metals.

0

50

100

150

200

250

Radi

us /

pm

F–

F

CIBr

I

CI–Br–

I–

Group 7

Figure 4.11 A comparison of size between halogens and their ions.

0

50

100

150

200

250

300

Radi

us /

pm

Na+

Mg2+

AI3+

Si4+

Si 4–

P 3 –

S 2 –

CI –

Period 3

Figure 4.12 Variation of ionic radius of positive and negative ions across period 3.

Na+

11+

Mg2+

12+

Figure 4.13 Mg2+ is smaller than Na+.


Melting pointAll the trends we have investigated so far (electronegativity, atomic radius, ionisation energy) have varied in the same way down any group or across any period: that is, ionisation energy always decreases down a group and electronegativity always increases across a period. The variation of melting point of the elements down a group changes from group to group and depends on the type of bonding in the element.

15+

P3–

16+

S2–

15+

P3–

16+

S2–

Test yourself1 Give the names of the following elements: a the element in period 3 and group 4 b the element in period 5 and group 6 c the element in the same group as sulfur but in

period 6 d a halogen in period 5 e an element in the same period as potassium that

has � ve outer shell electrons

2 State whether the following properties increase or decrease across a period:

a electronegativity b atomic radius

3 Arrange the following in order of increasing radius (smallest � rst):

a Ba Mg Sr Ca b O2− Na+ F−

c Na Na+ K Al3+

d S Cl I− Cl− S2−

4 Are the following true or false? a A germanium atom is smaller than a silicon atom,

but silicon has a higher � rst ionisation energy. b Selenium has a higher � rst ionisation energy

and electronegativity than sulfur. c Antimony has a higher � rst ionisation energy

and electronegativity than tin. d Cl− is bigger than Cl, but Se2− is smaller than Se. e Iodine has a higher electronegativity than

tellurium but a lower electronegativity than bromine.

Self-test 1

Test yourselfTest yourself1 Give the names of the following elements:

Self-test 1

1

Figure 4.14 S2− is smaller than P3−.

Now let us consider P3− and S2−: both ions have the same number of electrons. S2− has the higher nuclear charge, and therefore, as the amount of electron–electron repulsion is the same in both ions, the electrons are pulled in more strongly in S2− (Figure 4.14).

152

0

200

150

100

50

Mel

ting

poin

t / °C

Li

Na

K

Rb

Cs

Group 1

Variation in melting point down group 1The variation in melting point down group 1 (the alkali metals):

Melting point decreases down group 1.

+ +

+ +

– –

smaller distancebetween nucleusand delocalisedelectrons

Li Li

Li Li

Rb Rb

Rb Rb

++

++

Figure 4.15 The delocalised electrons are attracted more strongly in lithium than in rubidium.

The bonding in all these elements is metallic. The solid is held together by an electrostatic attraction between the positive ions in the lattice and the delocalised electrons (see page 137).

The attraction for the delocalised electrons is actually due to the nucleus of the positive ion, and therefore as the ion gets larger as we go down the group, the nucleus becomes further from the delocalised electrons and the attraction becomes weaker (Figure 4.15). This means that less energy is required to break apart the lattice as we go down group 1.

Variation in melting point down group 7The variation of melting point of halogens down group 7 (the halogens):

–250

–200

–150

–100

–50

0

50

100

150

Mel

ting

poin

t /

°C

F2

Cl2

l2

Br2Group 7 halogens

Melting point increases down group 7.

Liquid sodium is used as a coolant in some nuclear reactors.

As the relative molecular masses of the X2 halogen molecules increase, the van der Waals’ forces between molecules get stronger. This means that more energy must be supplied to separate the molecules from each other.

Variation of melting point across period 3In order to understand the variation in melting point across period 3, we must look at how the structure and bonding in the elements changes across the period (see Table 4.1 and Figure 4.16). The melting point


–400

–200

0

200

400

600

800

1000

1200

1400

1600

Mel

ting

poin

t / °C

Na

Mg

Si

AI

P

S

ArCIPeriod 3

Na Mg Al Si P S Cl Ar

Melting point (°C) 98 649 660 1410 44 119 −101 −189

Boiling point (°C) 883 1090 2467 2355 280 445 −34 −186

Structure metallic giant covalent covalent molecular atomic

Species present Na+ ions Mg2+ ions Al3+ ions Si atoms P4 molecules S8 molecules Cl2 molecules Ar atoms

Table 4.1 Variation in melting point across period 3.

Figure 4.16 Variation in melting point across period 3.

+ + + + +

+ + + + +

+ + + + +

+ + + + +

+ +

Na Mg

+ + +

2+ 2+ 2+ 2+ 2+

2+ 2+ 2+ 2+ 2+

2+ 2+ 2+ 2+ 2+

2+ 2+ 2+ 2+ 2+

2+ 2+ 2+ 2+ 2+

Figure 4.17 Metallic bonding in sodium and magnesium.

increases from sodium to magnesium to aluminium, which all have metallic bonding. There is then a substantial increase to the melting point of silicon, which has a giant covalent (macromolecular) structure. After silicon there is a sharp decrease in melting point as the structure and bonding change from giant covalent to covalent molecular.

The reason the melting point increases from sodium to magnesium to aluminium can be understood by comparing the metallic bonding in sodium and magnesium (Figure 4.17). There are several reasons why

154

magnesium has a higher melting point than sodium. The � rst of these is that magnesium forms a 2+ ion, whereas sodium forms a 1+ ion. This means that the electrostatic attraction between the ions and the delocalised electrons is stronger in magnesium. Also, there are two delocalised electrons per atom in magnesium, so there will be a greater number of electrostatic attractions between the ions and the delocalised electrons. Thirdly, the Mg2+ ion (65 pm) is smaller than the Na+ ion (98 pm), and therefore the delocalised electrons are closer to the nucleus of the positive ion in magnesium and more strongly attracted.

Silicon has a very high melting point because it has a giant covalent structure, and covalent bonds must be broken when it is melted. Covalent bonds are very strong, and a lot of energy is required to break them.

There are various allotropes (structural modi� cations) of phosphorus. The data given in Table 4.1 are for white phosphorus, which is the most common form of phosphorus. It consists of P4 tetrahedra with van der Waals’ forces between them. These P4 tetrahedra also exist in the liquid state, and therefore only van der Waals’ forces are broken when phosphorus is melted. Van der Waals’ forces are weak and little energy is required to break them. The melting point of phosphorus is therefore much lower than that of silicon, because, when phosphorus is melted, only weak van der Waals’ forces are broken rather than covalent bonds.

Sulfur, like phosphorus, has a covalent molecular structure, with van der Waals’ forces between molecules. A sulfur molecule, however, consists of eight atoms and, as S8 has a signi� cantly higher relative molecular mass than P4, the van der Waals’ forces are stronger between S8 molecules than between P4 molecules. More energy is thus required to break the van der Waals’ forces between S8 molecules than between P4 molecules, and sulfur has a higher melting point than phosphorus.

Going fom sulfur to argon, the realtive mass of the molecule (atom in the case of argon) decreases. Sulfur exists as S8 molecules (Mr 256.48), chlorine as Cl2 molecules (Mr 70.90) and argon as atoms (Ar 39.95). Thus, the decrease in melting point from sulfur to argon can be attributed to a decrease in van der Waals’ forces as the mass of the particles decreases.

4.3 Chemical properties of elements in group 1 and group 7

Chemical properties of elements in the same groupThe reactions of an atom are determined by the number of electrons in the outer shell (highest main energy level), and as elements in the same group in the periodic table have the same number of electrons in their outer shell, they react in basically the same way.

covalentbond

Si atom

covalentbond

P atom

covalentbond

S atom

Learning objectives

• Understand that elements in the same group have similar chemical properties

• Describe some reactions of elements in group 1 and group 7


Reactions of the elements in group 1The elements in group 1 are all reactive metals that react readily with, among other things, oxygen, water and halogens. The atoms all have one electron in their outer shell, and reactions virtually all involve the loss of this outer shell electron to form the positive ion, M+. The reactions become more vigorous as we descend through the group, because the ionisation energy decreases as the size of the atom increases. This means that, for example, Cs loses its outer electron to form a positive ion much more easily than Na and will react more vigorously.

Reaction with oxygenThe alkali metals react vigorously with oxygen and all tarnish rapidly in air. The general equation for the reaction is:

4M(s) + O2(g) → 2M2O(s)

Reaction with waterThe alkali metals react rapidly with water. The general equation for the reaction is:

2M(s) + 2H2O(l) → 2MOH(aq) + H2(g)

An alkaline solution is formed. The alkali metal hydroxides are strong bases and ionise completely in aqueous solution (page 319).

The reaction with water becomes more vigorous as we descend down the group: thus, sodium melts into a ball, � zzes rapidly and moves around on the surface of the water, potassium bursts into � ames (lilac) and caesium explodes as soon as it comes into contact with water.

Reactions of the elements in group 7The atoms of the elements in group 7 all have seven electrons in their outer shell and react either by gaining an electron to form the X− ion or by forming covalent compounds. The reactivity decreases down the group, and � uorine is the most reactive element known, reacting directly with virtually every other element in the periodic table. The variation in reactivity of the halogens cannot be as easily explained as for the alkali metals. The very high reactivity of � uorine can be explained in terms of an exceptionally weak F–F bond and the strength of bonds it forms to other atoms. The reactivity in terms of the formation of X− ions can be related to a decrease in electron a� nity (energy released when an electron is added to a neutral atom) down the group as the electron is added to a shell further away from the nucleus, but this is only part of the story, and several factors must be considered when explaining the reactivity of the halogens.

The halogens all react with the alkali metals to form salts. The general equation is:

2M(s) + X2(g) → 2MX(s)

The salts formed are all white/colourless, fairly typical ionic compounds. They contain M+ and X− ions. All alkali metal chlorides, bromides and iodides are soluble in water and form colourless, neutral solutions.

M2O is a basic oxide that will dissolve in water to form an alkaline solution, containing M+ and OH− ions.

Li, Na and K are all less dense than water.

How vigorous the reaction is depends on the particular halogen and alkali metal used; the most vigorous reaction occurs between � uorine and caesium, and the least vigorous reaction between lithium and iodine.

Chlorine is produced by the electrolysis of brine (see Option C). Worldwide annual production is about 60 million tons. Chlorine and its compounds are involved in the production of about 90% of the most important pharmaceuticals. Its biggest single use is in the production of PVC.

156

Displacement reactions of halogensThese are reactions between a solution of a halogen and a solution containing halide ions and are discussed in more detail on page 388. A small amount of a solution of a halogen is added to a small amount of a solution containing a halide ion, and any colour changes are observed (see Table 4.2). Potassium chloride, bromide and iodide solutions are all colourless. The colours of chlorine, bromine and iodine solutions are shown in Figure 4.18.

KCl(aq) KBr(aq) KI(aq)

Cl2(aq) no reaction orange solution dark red/brown solution

Br2(aq) no reaction no reaction dark red/brown solution

I2(aq) no reaction no reaction no reaction

Table 4.2 Results of reactions between halogen solutions and solutions containing halide ions.

The reactions that occur are:

Cl2(aq) + 2KBr(aq) → 2KCl(aq) + Br2(aq)Ionic equation: Cl2(aq) + 2Br−(aq) → 2Cl−(aq) + Br2(aq)

Cl2(aq) + 2KI(aq) → 2KCl(aq) + I2(aq)Ionic equation: Cl2(aq) + 2I−(aq) → 2Cl−(aq) + I2(aq)

Br2(aq) + 2KI(aq) → 2KBr(aq) + I2(aq)Ionic equation: Br2(aq) + 2I−(aq) → 2Br−(aq) + I2(aq)

The more reactive halogen displaces the halide ion of the less reactive halogen from solution. Thus, chlorine displaces bromide ions and iodide ions from solution, and bromine displaces iodide ions from solution.

4.4 Properties of the oxides of period 3 elements

Oxides of period 3 elementsTable 4.3 shows the formulas, structure and bonding of the oxides of period 3 elements.

Sodium oxide, magnesium oxide and aluminium oxideThese all have giant ionic structures. The ions are held in the lattice structure by strong electrostatic forces, and therefore they all have high melting points. MgO and Al2O3, which consist of highly charged ions, both have melting points of over 2000 °C. When the solids are melted ionic liquids are formed. The ions are free to move around in the liquids, and they all conduct electricity.

CI2(aq) Br2(aq) I2(aq)

Figure 4.18 Chlorine solution is pale yellow–green, bromine solution is orange, and iodine solution is red–brown.

Orange colour, due to the production of bromine.

Red-brown colour, due to the production of iodine.

Learning objectives

• Relate trends in the properties of the oxides of period 3 elements to their structure and bonding

• Describe the reactions of period 3 oxides with water


Covalent oxidesThe structures of the covalent oxides of period 3 elements are shown in Table 4.4. Substances with van der Waals’ forces/dipole–dipole interactions between molecules have low melting points, as these forces are weak and little energy is required to break them.

Notes on the structuresP4O6 is a covalent molecular liquid at room temperature that freezes at 24 °C. The structure is based on a tetrahedron of phosphorus atoms with an oxygen atom bridging each edge. There is no direct bonding between the P atoms. P4O10 has a structure that is basically the same as that of P4O6 except that there are four additional O atoms doubly bonded to each P atom.

In SO2, there are three negative charge centres around the S atom, and so the molecule is based on a trigonal planar shape (see Figure 4.19). As one of the negative charge centres is a lone pair, the shape is bent. O is more electronegative than S, and the molecule is polar. (It is also possible to draw an alternative Lewis structure for SO2 in which the octet on the S is not expanded – this is discussed on page 97).

Sulfur(VI) oxide (also called sulfur trioxide) is a colourless liquid at room temperature. The structure of the liquid is complex, and SO3 molecules trimerise (three molecules join together) to a certain extent to form S3O9. The structure of the molecule present in the gaseous state is shown in the diagram in Table 4.4 (page 158).

Examiner’s tipOnly part of this section is required by Standard Level students – if you are studying Standard Level Chemistry, you should understand that the bonding in the oxides changes from ionic to covalent across the period. You should not have to recall the structures of the oxides.

OOS

Figure 4.19 The Lewis structure for SO2.

Na Mg Al Si P S Cl

Name of oxide sodium oxide

magnesium oxide

aluminium oxide

silicon(IV) oxide

phosphorus(III) oxide

sulfur(IV) oxide chlorine(I) oxide

phosphorus(V) oxide

sulfur(VI) oxide chlorine(VII) oxide

Formula of oxide Na2O MgO Al2O3 SiO2 P4O6 SO2 Cl2O

P4O10 SO3 Cl2O7

Physical state at 25 °C solid

liquid gas gas

solid liquid liquid

Bonding in oxide ionic covalent

Structure giant simple molecular

Species present in liquid state

Na+ and O2− ions

Mg2+ and O2− ions

Al3+ and O2− ions

Si and O atoms

P4O6 molecules SO2 molecules Cl2O molecules

P4O10 molecules SO3 molecules Cl2O7 molecules

Electrical conductivity when molten

good none

liquid oxides conduct electricity as mobile ions present in liquid state

do not conduct electricity in liquid state as no ions present.

Table 4.3 Properties of the oxides of the period 3 elements.

158

Although SO3 is non-polar and SO2 is polar, SO3 has a higher relative molecular mass than SO2, meaning that the van der Waals’ forces between molecules are stronger and that SO3 has a higher melting point than SO2.

Cl2O7 is a colourless liquid at room temperature. Its structure involves two ClO4 tetrahedra linked by a bridging oxygen atom.

Reactions of period 3 oxides with waterA summary of the structure and bonding of the period 3 oxides and their reactions with water is shown in Table 4.5.

In general, metallic oxides are basic and non-metallic oxides are acidic.

Silicon Phosphorus Sulfur Chlorine

Bonding

in oxide

giant covalent covalent molecular covalent molecular covalent molecular

Structure giant simple molecular

SiO2 P4O6 SO2 Cl2O

covalent bonds broken when melted – high melting point

van der Waals’ forces between molecules – low melting point

dipole–dipole interactions and van der Waals’ forces between molecules – low melting point

Structure

P4O10 SO3 Cl2O7

van der Waals’ forces between molecules – low melting point; higher melting point than P4O6 due to higher Mr

van der Waals’ forces between molecules – low melting point; higher melting point than SO2 due to higher Mr

dipole–dipole interactions and van der Waals’ forces between molecules – low melting point; higher melting point than Cl2O due to higher Mr

Table 4.4 The structure of covalent oxides of the period 3 elements.

covalentbond

O atom

Si atom

covalentbond

O atom

P atom

S atom Cl2O

Cl atom

covalentbond

O atom

P atom

S atom

CI atom

Extension

Although there are four pairs of electrons around the central, bridging O atom in Cl2O7, the bond angle is 119°. Can you provide an explanation for this?


Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine

Formula of oxide Na2O MgO Al2O3 SiO2 P4O6 SO2 Cl2O

P4O10 SO3 Cl2O7

Nature of element metal non-metal

Bonding in oxide ionic giant covalent covalent molecular

Nature of oxide basic amphoteric acidic

Reaction with water soluble and reacts sparingly soluble, some reaction

insoluble soluble and reacts

Solution formed alkaline slightly alkaline – acidic

Table 4.5 Bonding and reaction with water in the period 3 oxides.

A basic oxide is one that will react with an acid to form a salt and, if soluble in water, will produce an alkaline solution. Sodium oxide reacts with water to form sodium hydroxide, according to the equation:

Na2O(s) + H2O(l) → 2NaOH(aq)

Sodium oxide reacts with acids such as sulfuric acid to form salts:

Na2O(s) + H2SO4(aq) → Na2SO4(aq) + H2O(l)

Magnesium oxide, because of the high charges on the ions, is not very soluble in water, but it does react to a small extent to form a solution of magnesium hydroxide, which is alkaline:

MgO(s) + H2O(l) → Mg(OH)2(aq)

Aluminium is right on the dividing line between metals and non-metals and forms an amphoteric oxide (some of the properties of a basic oxide and some of an acidic oxide). Aluminium is here exhibiting properties that are between that of a metal (basic oxide) and a non-metal (acidic oxide).

Aluminium oxide does not react with water, but it displays amphoteric behaviour in that it reacts with acids and bases to form salts.

Reaction with acids:

Al2O3 + 6H+ → 2Al3+ + 3H2O

Reaction with alkalis/bases:

Al2O3 + 2OH− + 3H2O → 2Al(OH)4−

The remaining oxides in Table 4.5 are all acidic oxides. An acidic oxide is one that reacts with bases/alkalis to form a salt and, if soluble in water, will produce an acidic solution.

Examiner’s tipThe reactions highlighted like this one must be learnt for examinations.

Amphoteric oxide – reacts with both acids and bases.

160

P4O6 (phosphorus(III) oxide) and P4O10 (phosphorus(V) oxide) form phosphoric(III) and phosphoric(V) acid, respectively, when they react with water:

P4O6(s) + 6H2O(l) → 4H3PO3(aq)

P4O10(s) + 6H2O(l) → 4H3PO4(aq)

SO2 (sulfur(IV) oxide) and SO3 (sulfur(VI) oxide) form sulfuric(IV) and sulfuric(VI) acid, respectively, when they react with water:

SO2(g) + H2O(l) → H2SO3(aq)

SO3(g) + H2O(l) → H2SO4(aq)

Cl2O (chlorine(I) oxide) and Cl2O7 (chlorine(VII) oxide) form acidic solutions (chloric (I) acid and chloric (VII) acid, respectively) when they react with water:

Cl2O(g) + H2O(l) → 2HClO(aq)

Cl2O7(l) + H2O(l) → 2HClO4(aq)

Test yourself5 Write balanced equations for the following

reactions: a rubidium with water b potassium with bromine c chlorine solution with potassium bromide

solution d sodium oxide with water e sulfur(VI) oxide with water

6 State the type of structure and bonding in each of the following:

a SO2 b SiO2 c Na2O d Al2O3

7 Arrange the following in order of increasing melting point (lowest � rst):

a Cl2 Na F2 K b Si Mg Cl2 Ar Ne c Na2O P4O10 O2

8 State whether an acidic or alkaline solution will be formed when each of the following is dissolved in/reacted with water:

a SO3

b MgO c Na

Self-test 2

Test yourselfTest yourself5 Write balanced equations for the following

Self-test 2

5 Write balanced equations for the following

Non-metal oxides such as SO2 are produced in various industrial processes and when coal is burnt. This can be responsible for acid rain, which can, among other things, kill � sh in lakes and trees in forests.

Nitrogen oxides (NOx) may be formed in internal combustion engines, and these are involved in the formation of photochemical smog in cities.

Figure 4.20 A photochemical smog over Hong Kong.

Phosphoric(V) acid is an ingredient of Coca-Cola®.


4.5 Properties of the chlorides of period 3 elements

Chlorides of period 3Table 4.6 shows some data for some of the chlorides of period 3 elements.Sodium and magnesium chloride are both high-melting-point ionic solids. They consist of a giant lattice of positive and negative ions (Figure 4.21). The oppositely charged ions attract each with strong electrostatic forces, and therefore a lot of energy is required to separate the ions. The bonding in these compounds is considered on pages 84–90.

When NaCl or MgCl2 are melted, the ions become free to move around, and so the molten salts conduct electricity.

HL

Na Mg Al Si P Cl

Formula of chloride NaCl MgCl2 ‘AlCl3’ = Al2Cl6 SiCl4 PCl3, PCl5 Cl2

Bonding in chloride ionic ionic covalent molecular*

covalent molecular

covalent molecular**

covalent molecular

State at 25 °C solid solid solid liquid PCl3 liquid PCl5 solid

gas

Electrical conductivity

of molten chloride

good good none none none none

Species present in

liquid state

Na+ and Cl− ions

Mg2+ and Cl− ions

Al2Cl6 molecules

SiCl4 molecules PCl3 / PCl5 molecules

Cl2 molecules

Table 4.6 Properties of the chlorides of the period 3 elements.

CI– Na+

The bonding in aluminium chloride is complicated, and it undergoes a change in structure and bonding as it changes state. For simplicity, we will assume that the bonding is covalent molecular in all states. The Lewis structure for AlCl3 shows that the Al only has six electrons in its outer shell. Figure 4.21 The ionic lattice in sodium

chloride.

Extension

* Actually bonding in the solid state is more ionic with six coordinate Al present. Covalent molecular (Al2Cl6) is present in liquid and gaseous states.

** PCl5 exists as [PCl6]−[PCl4]+ in the solid state but is covalent molecular in the liquid state.

CIAICICI

CIAICICI

CIAI CICI

CIAI CI

AI CICI

CICI

dativecovalentbond

dativecovalentbond

dativecovalentbond

dativecovalentbond

In order to generate an octet, two AlCl3 molecules come together to form a dimer. A dative covalent bond is formed from a lone pair on a Cl atom on one AlCl3 unit to an Al on the other AlCl3 unit.

Learning objectives

• Relate trends in the properties of the chlorides of period 3 elements to their structure and bonding

• Describe the reactions of period 3 chlorides with water

162

With four pairs of electrons around each Al atom, the shape about each Al atom is approximately tetrahedral, and the Al2Cl6 molecule is non-planar.

There are van der Waals’ forces between the Al2Cl6 molecules, and as these are relatively weak, little energy is required to break them. This means that Al2Cl6 has a low melting point and boiling point. Because only Al2Cl6 molecules (and no ions) are present in liquid aluminium chloride it does not conduct electricity.

SiCl4 and PCl3 are covalent molecular liquids. They consist of individual SiCl4 or PCl3 molecules. SiCl4 is tetrahedral, whereas PCl3 has a trigonal pyramidal (based on tetrahedral) shape (see Figure 4.22).

In both cases there are van der Waals’ forces between the molecules and, as little energy is required to break these, they have low melting points and boiling points – both are liquids at room temperature. As there are no ions present in liquid silicon chloride or liquid phosphorus(III) chloride, but only SiCl4 or PCl3 molecules, neither substance conducts electricity.

The bonding in PCl5 is a little more complex, as it actually consists of octahedral [PCl6]− and tetrahedral [PCl4]+ ions in the solid state. This ionic bonding leads to PCl5 being a solid at room temperature (although the electrostatic forces between these large ions are not very strong and the melting point of PCl5 is only 167 °C). When PCl5 melts it forms a covalent molecular liquid containing trigonal bipyramidal molecules. The forces between these molecules are van der Waals’ forces. As PCl5 has a higher relative molecular mass than PCl3, the van der Waals’ forces between molecules would be expected to be stronger between the PCl5 molecules, leading to PCl5 having a higher boiling point than PCl3.

Only PCl5 molecules are present in the liquid state, and therefore, as no ions are present, liquid PCl5 does not conduct electricity.

Cl2 is a covalent molecular gas consisting of diatomic molecules. There are only weak van der Waals’ forces between the molecules in liquid chlorine and, as Cl2 has a relatively low relative molecular mass, these are weak enough to be broken at temperatures below room temperature. This means that Cl2 is a gas at room temperature. There are no ions present in liquid chlorine, and so it does not conduct electricity.

Reactions of period 3 chlorides with waterTable 4.7 summarises the reactions of period 3 chlorides with water.

CI CICI

a b

PCI CICI

Si

CI

Figure 4.22 (a) The SiCl4 molecule. (b) The PCl3 molecule.

CI

CICI

CI

CI

P

HL

Sodium Magnesium Aluminium Silicon Phosphorus Chlorine

Formula of chloride NaCl MgCl2 ‘AlCl3’ = Al2Cl6 SiCl4 PCl3, PCl5 Cl2

Reaction with water dissolves, no reaction

dissolves, very slight reaction

fairly vigorous reaction with water; fumes in moist air

vigorous reaction

vigorous reaction

dissolves and reacts slightly

pH (aqueous solution)* 7 6.2 3 1 1 2

Acidity neutral slightly acidic acidic acidic acidic acidic

* pH depends on concentration – these are approximate values for a 0.1 mol dm−3 solution.

Table 4.7 Reaction with water in the period 3 chlorides.

CI

CI

CI

CI

CI

CIAI AI


Sodium chloride dissolves in water. The ions separate from each other and become hydrated:

excess waterNaCl(s) ⎯⎯⎯⎯→ Na+(aq) + Cl−(aq)

Magnesium chloride also dissolves in water (with the liberation of a great deal of heat), but because of the higher charge density (higher charge : radius ratio) of Mg2+ ions compared with Na+ ions there is some polarisation of water, which causes the release of some H+ ions from water. This can be represented by the equation:

[Mg(H2O)6]2+(aq) [Mg(H2O)5(OH)]+(aq) + H+(aq)

The charge on the Mg2+ ion is not very high, however, and the position of equilibrium lies mostly to the left.

Anhydrous aluminium chloride fumes in moist air because of the production of hydrogen chloride gas. This reaction can be represented by an equation such as:

AlCl3(s) + 3H2O(l) Al(OH)3(s) + 3HCl(g)

With excess water, the aluminium chloride dissolves according to the equation:

AlCl3(s) + 6H2O(l) → [Al(H2O)6]3+(aq) + 3Cl−(aq)

The acidity of the solution is a result of the ionisation of the aqueous Al3+ ion:

[Al(H2O)6]3+(aq) [Al(H2O)5(OH)]2+(aq) + H+(aq)

This occurs to a much greater extent than with Mg2+, as the Al3+ ion is smaller and more highly charged than the Mg2+ ion and causes greater polarisation of water molecules.

Silicon chloride reacts very vigorously with water, with the production of a great deal of heat and fumes of hydrogen chloride gas. The SiCl4 is hydrolysed completely by water, and the reaction can be represented by an equation such as:

SiCl4(l) + 4H2O(l) → Si(OH)4(s) + 4HCl(aq) hydrochloric acid

The acidity of the solution is due to the production of HCl in the reaction, which dissolves in the water to produce hydrochloric acid.

Phosphorus(III) chloride reacts very vigorously in water, with the production of heat and fumes of hydrogen chloride gas. Phosphoric(III) acid and hydrochloric acid are formed:

PCl3(l) + 3H2O(l) → H3PO3(aq) + 3HCl(aq) phosphoric(III) hydrochloric acid acid

This is discussed in more detail on pages 357 and 358.

HL

Extension

The reaction is actually more complicated than this, and Si(OH)4 undergoes condensation polymerisation to form SiO2.xH2O.

164

HL Phosphorus(V) chloride reacts very vigorously with water, with the production of heat and fumes of hydrogen chloride gas. Phosphoric(V) acid and hydrochloric acid are formed:

PCl5(l) + 4H2O(l) → H3PO4(aq) + 5HCl(aq) phosphoric(V) hydrochloric acid acid

Chlorine dissolves in water and reacts to a certain extent. The relevant reactions are:

excess waterCl2(g) ⎯⎯⎯→ Cl2(aq)

Cl2(aq) + H2O(l) HCl(aq) + HOCl(aq) hydrochloric acid chloric(I) acid

The solution is acidic mainly as a result of the production of hydrochloric acid. Chloric(I) acid is a very weak acid (pKa = 7.43).

26Fe

27Co

28Ni

29Cu

30Zn

25Mn

24Cr

23V

22Ti

21Sc

Fe Co Ni Cu ZnMnCrVTi[Ar]4s23d1 [Ar]4s23d2 [Ar]4s23d3 [Ar]4s13d5 [Ar]4s23d5 [Ar]4s23d6 [Ar]4s23d7 [Ar]4s23d8 [Ar]4s13d10 [Ar]4s23d10

Sc

They are called d-block elements because the subshell being � lled across this series is the 3d subshell. The electronic con� gurations range from [Ar]4s23d1 for scandium to [Ar]4s23d10 for zinc.

There are also two other rows of d-block elements.

Examiner’s tipRemember that Cr and Cu have slightly di� erent electronic con� gurations.

Extension

Cl2 dissolves better in alkaline solution, as the equilibrium is shifted to the right.

9 Arrange the following in order of increasing melting point (lowest � rst):

a Cl2 SiCl4 PCl5 NaCl b PCl3 SiCl4 MgCl2 AlCl3 10 State whether each of the following is true or

false. a Liquid sodium chloride conducts electricity

but liquid phosphorus(III) chloride does not.

b SiCl4 and NaCl both have giant structures. c SiCl4 and PCl5 are both liquids at room

temperature and pressure. d PCl3 and PCl5 both react with water to form

H3PO4. e MgCl2 reacts with water to form a neutral

solution, but SiCl4 reacts to form an acidic solution.

Test yourself

4.6 The transition elements

The transition elements (d block)The � rst row d-block elements are:

Learning objectives

• Describe the characteristic properties of transition metals

• Explain why transition metals have variable oxidation states

• Explain the formation and describe the shape of complex ions

• Explain why transition metal complex ions are coloured

• Describe some uses of transition metals and their compounds as catalysts


HL

Fe Co Ni Cu ZnMn

Transition elements

CrVTiSc

The transition elements can be de� ned as di� erent from the d-block elements, and the de� nition we will use here is:

a transition element is an element that forms at least one stable oxidation state (other than 0) with a partially � lled d subshell

According to this de� nition, scandium and zinc are not counted as transition elements, as the only ion formed by Sc is the 3+ ion, with electronic con� guration 1s22s22p63s23p6 (no d electrons), and that formed by Zn is the 2+ ion, with electronic con� guration 1s22s22p63s23p63d10 (full d subshell).

Scandium and zinc do not exhibit some of the characteristic properties of transition metals detailed below (e.g. they do not form coloured compounds).

Properties of the transition elementsWe have already studied the variation in properties of a set of eight elements across the periodic table when we looked at the properties of period 3 elements. The transition elements also form a set of eight elements across the periodic table, but these are much more similar to each other than the elements across period 3. For instance, they are all metals rather than showing a change from metal to non-metal.

The variation in � rst ionisation energy and atomic radius of the transition elements and period 3 elements are compared in Figures 4.23 and 4.24. It can be seen that the variation of ionisation energy and atomic radius across the series of the transition elements is much smaller than across period 3.

These are often called transition metals rather than transition elements.

0

200

400

600

800

1000

1200

1400

1600

Firs

t io

nisa

tion

ener

gy /

kJ m

ol –1

Ti

Na

Mg

V

Cr

Si

MnFe

P

S

Co NiCu

Ar

CI

AI

Period 3 / transition elements

Period 3

Key

Transition elements

Figure 4.23 A comparison of the variation of fi rst ionisation energy across period 3 with that across the transition metal series.

166

200

180

160

140

120

100

80

Ato

mic

rad

ius /

pm

Ti

Na

Mg

V CrSi

Mn Fe

P S

Co Ni Cu

CI

AI

Period 3 / transition elements

Period 3

Key

Transition elements

Figure 4.24 A comparison of the variation of atomic radius across period 3 with that across the transition metal series.

Examiner’s tipThe last four properties are most important for examinations.

The 4s electrons are always removed before the 3d electrons when an ion is formed.

HL

Because of their similarity it is possible to draw up a list of characteristic properties of transition elements:• Transition elements are all typical metals, i.e. they have high melting

points and densities.• Transition elements can exhibit more than one oxidation number in

compounds/complexes.• Transition elements form complex ions.• Transition elements usually form coloured compounds/complexes.• Transition elements and their compounds/complexes can act as catalysts

in many reactions.

Ionisation of transition elementsTransition elements form positive ions. The electronic con� gurations of some transition metal ions are shown in Table 4.8.

Element Electronic confi guration Ion Electronic confi guration

Cr [Ar]4s13d5 Cr2+ [Ar]3d4

Cr3+ [Ar]3d3

Mn [Ar]4s23d5 Mn2+ [Ar]3d5

Fe [Ar]4s23d6 Fe2+ [Ar]3d6

Fe3+ [Ar]3d5

Co [Ar]4s23d7 Co2+ [Ar]3d7

Cu [Ar]4s13d10 Cu+ [Ar]3d10

Cu2+ [Ar]3d9

Table 4.8 Electronic confi gurations of transition metals and their ions.


Variable oxidation numbersThe positive oxidation numbers (oxidation states) exhibited by the transition elements are shown in Figure 4.25.

The greatest number of di� erent oxidation numbers and the highest oxidation numbers are found in the middle of the series. From Ti to Mn there is an increase in the total number of electrons in the 4s and 3d subshells, so the maximum oxidation number increases. Mn has the electronic con� guration [Ar]4s23d5 and therefore a maximum oxidation state of +7. Fe has eight electrons in the 4s and 3d subshells and would be expected to have a maximum oxidation state of +8, but the ionisation energy increases from left to right across the transition elements series and it becomes more di� cult to reach the highest oxidation numbers towards the right-hand side of the series. The chemistry of copper and nickel is, for the same reason, dominated by the lower oxidation numbers.

All transition metals show oxidation number +2. In most cases this is because they have two electrons in the 4s subshell, and removal of these generates an oxidation number of +2.

Examiner’s tipThe oxidation numbers highlighted in Figure 4.25 are mentioned speci� cally on the syllabus – you should remember these.

Mn

7Cr Fe

66V 6

555Ti 5 Co Ni Cu

4444 4 4 4 4

3333 3 3 3 3

2222 2 2 2 2

1111 1 1 1 1

0000 0 0 0 0

Figure 4.25 Oxidation numbers of transition metals in compounds. Not all oxidation numbers are common.

Oxidation number and oxidation state are the same thing.

HL

Why more than one oxidation number?The 4s and 3d subshells are close in energy, and there are no big jumps in the successive ionisation energies when the 4s and 3d electrons are removed; therefore the number of electrons lost will depend on a variety of factors, such as lattice enthalpy, ionisation energy and hydration enthalpy. Electrons are not removed to generate the nearest noble gas electronic con� guration. The graph in Figure 4.26 (overleaf ) shows a comparison of the � rst seven ionisation energies of Mg and Mn. It can be seen that there is a very large jump between the second and third ionisation energies of Mg, but there are no such jumps for Mn.

Oxidation numbers are discussed further on pages 376–378.

168

A ligand must possess a lone pair of electrons.

Mg

Mn

25 000

20 000

15 000

Ioni

satio

n en

ergy

/ k

J mol

–1

10 000

5 000

01 2 3 4 5 6 7

Number of ionisation energy8

Figure 4.26 Comparison of successive ionisation energies of Mg and Mn.

Complex ionsA complex ion consists of a central transition metal ion surrounded by ligands.

Ligands are negative ions or neutral molecules that use lone pairs of electrons to bond to a transition metal ion to form a complex ion. Dative covalent bonds (coordinate links) are formed between the ligand and the transition metal ion.

The structure of [Fe(H2O)6]2+ is shown in Figure 4.27. H2O is the ligand in this complex ion. The shape of this complex ion is octahedral and it is called the hexaaquairon(II) ion. Other ways of drawing this complex ion are shown in Figure 4.28.

A ligand is a Lewis base.

CI

OH

Fe2+

H dative covalentbond

ligand

a b

O O

O O

O

PCI CI

CI

Si

CI

HHH

HH

H

H

HH

H

Figure 4.27 A complex ion.O

H2+

Fe

H

O O

O O

OHH

H

HH

H

H

HH

HO

H2+

Fe

H

O O

O O

OHH

H

HH

H

H

HH

H

Figure 4.28 Alternative representations of the [Fe(H2O)6]2+ complex ion.

All transition elements, with the exception of Ti, form an octahedral complex ion with the formula [M(H2O)6]2+ in solution.

The oxidation number of a transition metal in a complex ionThe oxidation number of a transition metal in a complex ion may be worked out from the charges on the ligands. Ligands may be either neutral or negatively charged.

In [Fe(H2O)6]2+ all the ligands are neutral. The overall charge on the ion is just due to the Fe, so the oxidation number of Fe must be +2.

Neutral ligands 1− ligands

H2O Cl−

NH3 CN−

CO Br−

HL

Extension

There is a strong case for considering the bonding in a transition metal complex ion as having a signi� cant ionic component. Crystal � eld theory and ligand � eld theory consider the bonding from a more ionic point of view.


Transition metal complexes do not obey valence shell electron pair repulsion (VSEPR) theory rules, so although six coordinate complexes are virtually always octahedral, four coordinate complexes may be tetrahedral or square planar.

HLIn [Ni(CN)4]2− all the ligands have a 1− charge, so the total charge on all four ligands is 4−. The overall charge on the ion is 2−; therefore, the oxidation state of Ni must be +2 to cancel out 2− from the 4− charge.

Shapes of complex ionsThe shapes of some complex ions are shown in Table 4.9.

Complex ion

Formula [Fe(H2O)6]3+ [Fe(CN)6]3−

Shape octahedral octahedral

Oxidation number of metal +3 +3

Name hexaaquairon(III) ion hexacyanoferrate(III) ion

Complex ion

Formula [CuCl4]2− [Ni(CN)4]2−

Shape tetrahedral square planar

Oxidation number of metal +2 +2

Name tetrachlorocuprate(II) ion tetracyanonickellate(II) ion

Complex ion

Formula [Ag(NH3)2]+

Shape linear

Oxidation number of metal +1

Name diamminesilver(I) ion

Table 4.9 The shapes of some complex ions.

OH

3+

Fe

H

O O

O O

OHH

H

HH

H

H

HH

HC

N

N

N

N

3–

FeC

C

C

N

NC

C

CI2–

Cu

CI

CICI

C

N

N N

N

2–

Ni

C

C C

+AgH3N NH3

Some scientists believe that the bonding between a transition

metal and a ligand is purely ionic. All scientists have the same experimental data available to them – to what extent is scienti� c knowledge objective and to what extent is it a matter of interpretation and belief?

170

When iron(III) salts are dissolved in water, very little of the pale purple [Fe(H2O)6]3+(aq) ion is present. The [Fe(H2O)6]3+(aq) ion is unstable in aqueous solution, as the 3+ charge on the metal ion causes the water molecules to be strongly polarised so that the ion dissociates:

[Fe(H2O)6]3+(aq) [Fe(H2O)5(OH)]2+(aq) + H+

The [Fe(H2O)6]3+ ion can then only be prepared in strongly acidic solution.

If a high concentration of cyanide ions is added to a solution containing the Fe3+(aq) ion, the [Fe(CN)6]3− ion can be formed. We can represent the reaction as:

[Fe(H2O)6]3+(aq) + 6CN−(aq) [Fe(CN)6]3−(aq) + 6H2O(l) red/orange

If ammonia solution is added to a solution of silver nitrate, the colourless [Ag(NH3)2]+ ion is formed:

[Ag(H2O)2]+(aq) + 2NH3(aq) [Ag(NH3)2]+(aq) + 2H2O(l)

Formation of coloured complexesThe colours of some complex ions are shown in Table 4.10.

In a gaseous transition metal ion, all the 3d orbitals have the same energy – that is, they are degenerate. However, when the ion is surrounded by ligands in a complex ion, these d orbitals are split into two groups. In an octahedral complex ion there are two orbitals in the upper group and three orbitals in the lower group.

See page 357 – all 3+ ions in solution are acidic.

Le Chatelier’ principle – a strongly acidic solution means that the concentration of H+ ions is high and the position of equilibrium is shifted to the left-hand side.

Complex ion Colour

[Cu(H2O)6]2+ blue

[Cu(NH3)4(H2O)2]2+ deep blue/violet

[Fe(SCN)(H2O)5]2+ blood red

[Ni(H2O)6]2+ green

Table 4.10 The colours of some complex ions.

ML6

d orbitalsdegenerate

d orbitals split intotwo groups in a

complex ion

M2+(g)

ENERGY

2+

HL

Extension

The d orbitals in a complex ion are higher in energy than the d orbitals in an isolated ion.

As the HCl is added, the yellow [CuCl4]2− complex ion is formed. Therefore, as HCl is added the solution changes colour from blue to green (mixture of blue and yellow). According to Le Chatelier’s principle, the position of equilibrium shifts to the right as Cl− is added.

Formation of complex ionsComplex ions may undergo substitution reactions in which, for example, H2O ligands are replaced by other ligands.

Addition of concentrated hydrochloric acid to copper sulfate solution:

[Cu(H2O)6]2+(aq) + 4Cl−(aq) [CuCl4]2−(aq) + 6H2O(l) pale blue yellow


Energy in the form of a certain frequency of visible light can be absorbed to promote an electron from the lower set of orbitals to the higher set:

electron

light energy absorption of lightenergy causes anelectron to be

promoted to thehigher set of d orbitals

CuSO4(aq)

orange lightabsorbed

white light

orange lightmissing

when orange light is removedfrom white light it appears blue

When white light travels through copper sulfate solution, orange light is absorbed, promoting an electron from the lower set of d orbitals to the higher set. This means that the light coming out contains all the colours of the spectrum except orange and so appears blue, the complementary colour to orange.

White light is a mixture of all colours (frequencies) of visible light.

For a substance to appear coloured, certain frequencies of light in the visible region of the spectrum must be absorbed.

Formation of coloured complex ions/compounds requires the presence of a partially � lled d subshell.

HL Extension

The d orbitals are split in di� erent ways in di� erent-shaped complex ions.

Extension

Cr2O72− (orange), CrO4

2− (yellow) and MnO4

− (purple) are all very highly coloured, but they have no d electrons. They are coloured because of a di� erent mechanism from the one described here.

Figure 4.29 A Cu+ or Zn2+ ion has 10 3d electrons.

Let us consider the Sc3+ ion or the Ti4+ ion. These both have no electrons in the 3d subshell and so are colourless, as it is not possible to absorb energy to promote a 3d electron.

The Cu+ ion and the Zn2+ ion both have 10 3d electrons (Figure 4.29), and as there is no space in the upper set of orbitals it is not possible to promote an electron to the upper set of orbitals. No light in the visible region of the spectrum is absorbed, and these ions are colourless.

What do we mean when we say that a solution of copper sulfate is blue? Is blueness a property of copper sulfate solution, or is the blueness in our minds? What colour

would copper sulfate solution be in orange light? Or in the dark?

Catalytic abilityThe elements and their compounds/complexes can act as catalysts. For example, � nely divided iron is the catalyst in the Haber process for the production of ammonia:

N2(g) + 3H2(g) 2NH3(g)

172

C

H

H

C

R

H

C

R

H

C

H

H

This is discussed further on page 441.

HL Vanadium(V) oxide (V2O5) has been used as a catalyst for the conversion of sulfur(IV) oxide into sulfur(VI) oxide in the Contact process for production of sulfuric(VI) acid:

2SO2(g) + O2(g) 2SO3(g)

Manganese(IV) oxide (MnO2) is a commonly used catalyst in the decomposition of hydrogen peroxide:

2H2O2(aq) → O2(g) + 2H2O(l)

Finely divided nickel is used in the hydrogenation of alkenes to form alkanes, for example:

H2C=CH2(g) + H2(g) → H3C–CH3(g)

Cobalt is a constituent of vitamin B12 (cobalamin). The vitamin is not actually an enzyme, but rather a co-factor for enzymes. Most of the reactions catalysed by enzymes associated with vitamin B12 are rearrangement reactions of the type:

Palladium, platinum and rhodium are used in catalytic converters in cars to remove, as far as possible, harmful substances from the exhaust gases. Palladium and rhodium are second-row transition elements, and platinum is a third-row transition element. Reactions such as the following occur in a catalytic converter:

2CO + 2NO → 2CO2 + N2

2C8H18 + 25O2 → 16CO2 + 18H2O

2CO + O2 → 2CO2

Catalysts are important in the Haber and Contact processesA catalyst is used to speed up the rate of a reaction. As well as the obvious result of this – that the product is produced more quickly – catalysts are important in industrial processes for other reasons. For example, if a catalyst is used, the reaction may be run at a lower temperature, meaning less fuel must be burnt to provide energy. The key reactions in both the Haber process and the Contact process involve an exothermic reaction, and by using a catalyst, a lower temperature can be used, which improves the yield of the product (a catalyst does not itself a� ect the yield of a reaction, but using a lower temperature does). A catalyst allows the product to be produced more quickly and at a lower temperature and therefore improves the overall cost e� ectiveness of these processes.

The Haber and Contact processes are discussed on pages 296–300.

Catalytic converters are important in reducing pollution caused by the burning of fossil fuels in car engines.


Some chemistry of transition elementsHere we will look at the chemistry of some transition metal ions that ties together chemistry from several areas of the course.

IronThe most common oxidation states for iron are +2 and +3. Most iron(II) salts are oxidised slowly by the air to iron(III) salts. This can be seen if we consider the standard electrode potentials (see page 392) for the relevant reactions:

Fe3+(aq) + e− Fe2+(aq) E = +0.77 V

O2(g) + 4H+(aq) + 4e− 2H2O(l) E = +1.23 V

Oxygen is a strong enough oxidising agent to oxidise Fe2+ to Fe3+:

4Fe2+(aq) + O2(g) + 4H+(aq) → 4Fe3+(aq) + 2H2O(l) E cell = +0.46 V

HL

We have already seen that Fe3+(aq) is acidic because of the hydrolysis reaction:

[Fe(H2O)6]3+(aq) [Fe(H2O)5(OH)]2+(aq) + H+

As the hexa-aqua complex is not present, solutions of iron(III) are yellow and not violet.

Chromium(VI) and chromium(III)The most important chromium(VI) complexes are the yellow chromate(VI), [CrO4]2−, and the orange dichromate(VI), [Cr2O7]2−, ions. These two species exist in equilibrium in aqueous solution, the dichromate(VI) ion predominating in acidic solution:

2[CrO4]2− + 2H+ [Cr2O7]2− + H2O yellow orange

Addition of acid shifts the equilibrium position to the right in order to use up the H+ that has been added. If alkali is added, it reacts with the H+ ions, causing the position of equilibrium to shift to the left-hand side to replace the H+ ions that have been removed as far as possible.

The +6 oxidation state is strongly oxidising, and acidi� ed dichromate(VI) is used as an oxidising agent in organic chemistry, e.g. in the oxidation of primary alcohols to aldehydes and then carboxylic acids:

C2H5OH + [O] → CH3CHO + H2O

CH3CHO + [O] → CH3COOH

The � lm Erin Brockovich describes the true story of a woman’s � ght against a major corporation that was releasing carcinogenic chromium(VI) into the drinking water supply of a town in California, USA.

The value for the cell potential is positive, which indicates that the reaction is spontaneas. See page 392 for more details.

174

In the process the orange dichromate(VI) ion is reduced to the green Cr3+

ion:

Cr2O72−(aq) + 14H+(aq) + 6e− 2Cr3+(aq) + 7H2O(l) E = +1.33 V

The [Cr(H2O)6]3+ ion is actually violet. However, this does not exist in aqueous solution, as it is acidic (pKa = 4.1).

[Cr(H2O)6]3+ [Cr(H2O)5(OH)]2+ + H+

The +3 oxidation state is the most stable and common one for chromium.

Manganese(IV) and manganese(VII)The +7 oxidation state for manganese is typi� ed by the dark purple manganate(VII) ion (permanganate ion).

The MnO4− ion is tetrahedral and strongly oxidising in both acidic and

alkaline solution:

MnO4−(aq) + 8H+(aq) + 5e− Mn2+(aq) + 4H2O(l) E = +1.51 V

purple pale pink

In less strongly acidic or basic medium, manganese(IV) oxide is formed:

MnO4−(aq) + 4H+(aq) + 3e− MnO2(s) + 2H2O(l)

black

MnO2 is the only important compound with Mn in the +4 oxidation state. It is most commonly encountered as the catalyst for decomposition of hydrogen peroxide (see above).

MnO2 is strongly oxidising and is reduced to Mn2+, which is the most stable oxidation state for Mn:

MnO2(s) + 4HCl(aq) → MnCl2(aq) + Cl2(g) + 2H2O(l)

CopperThe most stable oxidation state of copper is +2, which is familiar as the pale blue [Cu(H2O)6]2+ ion in aqueous solutions of copper salts. Copper also, however, forms some copper(I) compounds and complexes.

In aqueous solution the copper(I) ion is unstable and disproportionates according to the equation:

2Cu+(aq) → Cu(s) + Cu2+(aq)

The favourability of this reaction can be seen from a consideration of the standard electrode potentials for the half reactions:

Cu2+(aq) + e− Cu+(aq) E = +0.15 V

Cu+(aq) + e− Cu(s) E = +0.52 V

Overall then:

2Cu+(aq) → Cu(s) + Cu2+(aq) E cell= +0.37 V

See Fe3+, above.

MnO2 is insoluble in water.

It will oxidise alcohols to aldehydes etc. It is a stronger oxidising agent than Cr2O7

2−.

Disproportionation – the same element is oxidised and reduced.

HL

Beryllium–copper alloys are used for springs and contacts in mobile phones.


Copper(I) iodide is formed when potassium iodide solution is added to a solution of copper(II) (e.g. copper(II) sulfate):

2Cu2+(aq) + 4I−(aq) → 2CuI(s) + I2(aq) white precipitate

This reaction is used in the determination of the concentration of a solution of Cu2+ ions. Excess potassium iodide is added to a sample of the Cu2+(aq), and the iodine released is titrated against a standard solution of sodium thiosulfate (Na2S2O3).

Most copper(I) compounds are white/colourless due to having a full 3d subshell, although an exception is the red copper(I) oxide, Cu2O.

HL

11 Give the full electronic con� guration of the following ions: a Ni2+ b Co3+ c V3+ d Mn4+

12 Give the oxidation state of the transition metal in each of the following complex ions

a [Ni(H2O)6]2+ b [Fe(CN)6]4−

c [MnCl4]2− d [Co(NH3)6]3+

e [Co(NH3)5Br]2+ f [Co(NH3)4Br2]+

g [FeO4]2− h Ni(CO)4 13 Which of the following compounds are likely to be coloured? TiF4 VF5 MnF3 CoF2 CuBr ZnCl2

Test yourself

Self-test 3

11 Test yourselfTest yourself

Self-test 3

11

Exam-style questions

1 Which of the following properties decrease in value down group 7?

A electronegativity and � rst ionisation energyB melting point and electronegativityC melting point and atomic radiusD ionic radius and � rst ionisation energy

2 Which of the following is a transition element?

A Te B Sb C Ba D V

3 Which of the following forms an alkaline solution when added to water?

A SO3 B Na2O C P4O10 D SiO2

176

4 Which of the following is true for two elements in the same group in the periodic table?

A they have the same physical propertiesB they have similar chemical propertiesC they have the same electronegativityD they have the same number of shells of electrons

5 A non-metallic element, X, forms a gaseous oxide with the formula X2O that reacts with water to form an acidic solution. The element X is most likely to be:

A Na B S C Cl D P

6 Which of the following chlorides produces a neutral solution when added to water?

A MgCl2 B NaCl C SiCl4 D PCl3

7 Which of the following chlorides has the highest melting point?

A MgCl2 B SiCl4 C PCl5 D PCl3

8 What is the oxidation state of chromium in K3[Cr(CN)6]?

A +6 B −6 C +3 D −3

9 Which of the following complex ions would be expected to be colourless?

A [Ni(H2O)6]2+ C [Cu(H2O)6]2+

B [Zn(H2O)6]2+ D [Co(H2O)6]2+

10 Which of the following is not a characteristic property of transition metals?

A they form complex ionsB they have full d subshellsC they exhibit more than one oxidation state in compoundsD they form coloured compounds

HL


11 a The atomic and ionic radii of some elements are given in the table:

Element Atomic radius / pm Ionic radius / pm

Na 186 98

Al 143 45

Cl 99 181

K 231 133

i Explain why the atomic radius of aluminium is smaller than that of sodium. [2] ii Explain why the ionic radius of aluminium is smaller than its atomic radius but the atomic

radius of chlorine is larger than its atomic radius. [4] iii Explain why the ionic radius of potassium is smaller than that of chlorine. [2]

b Explain the following in terms of structure and bonding: i Sodium has a higher melting point than potassium. [3] ii Silicon has a higher melting point than chlorine. [3]

c State the formula of phosphorus(V) oxide and write an equation for its reaction with water. [2]

12 a Explain the following in terms of structure and bonding: i Sodium chloride has a higher melting point than silicon tetrachloride. [3] ii Sodium chloride conducts electricity when molten but silicon tetrachloride does not. [2]

b Draw a diagram showing the bonding in aluminium chloride and write an equation to show its reaction with water. [3]

c What is the oxidation state of iron in the complex ion [FeCl6]4−? [1]

d i Explain why the complex ion [Cu(H2O)6]2+ is coloured. [3] ii When concentrated hydrochloric acid is added to an aqueous solution of copper sulfate, the

colour of the solution changes from blue to green. Write an equation for the reaction that occurs and describe the shape of the complex ion formed. [2]

e Give an example of a reaction, including an equation, in which a named � rst-row transition element acts as a catalyst. [2]

HL

178

HL

Summary

group number = number of outer shell electrons

Electronegativity and 1st ionisation energy decrease down a group.

Atomic radius and ionic radius increase down a group.

elements arranged in order of atomic number

arranged in groups – vertical columns

example period 3

oxides chlorides

Melting point increases and reactivity decreases down group.

examples

Negative ions are larger than the parent atom.

Positive ions are smaller than the parent atom.

group 1 – the alkali metals

Alkali metals react with halogens:

2M + X2 → 2MX

A more reactive halogen displaces a less reactive halogen, for example

Cl2(aq) + 2KBr(aq) → 2KCl(aq) + Br2(aq)

React vigorously with water:

2M + 2H2O → 2MOH + H2

To TRANSITION METALS on next page

periods – horizontal rows

Na 2O + H2O → 2NaOH

MgO + H2O → Mg(OH)2

P4O10 + 6H2O → 4H3PO4

SO3 + H2O → H2SO4

Electronegativity and 1st ionisation energy increase across a period.

Atomic radius decreases across a period.

Reactivity increases and melting point decreases down group.

group 7 – the halogens

THE PERIODIC

TABLE

acidic

SiO2 giant covalent

P4O6P4O10

covalent molecular

SO2SO3

Cl2OCl2O7

basicNa2O

ionicMgO

amphoteric Al2O3

ionicNaCl

MgCl2

AlCl3 + 3H2O → Al(OH) + 3HCl

AlCl3 + 6H2O → [Al(H2O)6]3+ + 6Cl−

SiCl4 + 4H2O → Si(OH)4 + 4H+ + 4Cl−

PCl3 + 3H2O → H3PO3 + 3H+ + 3Cl−

PCl5 + 4H2O → H3PO4 + 5H+ + 5Cl−

Cl2 + H2O → HOCl + H+ + Cl−

acidic solution

covalent molecular

AlCl3

SiCl4PCl3PCl 5Cl2


HL Can have more than 1 oxidation number in compounds or complexes.

Can form complex ions with ligands

examples

can act as catalysts in form of elements or compounds

Form coloured compounds and complexes.

Presence of ligands causes d orbitals to split into 2 groups with different energy levels.

A certain frequency of visible light is absorbed to promote an electron to a higher d orbital – the complementary colour is transmitted.

examples Cu: +1, +2

At least 1 oxidation state has a partially fi lled d subshell.

One of the oxidation states is always +2.

The 4s electrons are always removed before the 3d electrons.

Ligands are either negative ions or neutral molecules.

use lone pairs to bind to metal through dative covalent bonds

shapes

Fe: +2, +3

Mn: +2, +4, +7

Cr: +2, +3, +6

TRANSITION METALS

6 ligands 2 ligands4 ligands

octahedral lineartetrahedral square planar

e.g.[Fe(H2O)6]3+

e.g.[Ag(NH3)2]+

e.g.[CuCl4]2−

e.g.[Ni(CN)4]2−

iron, production of ammonia: N2 + 3H2 2NH3 (Haber process)

V2O5, production of sulfur dioxide: 2SO2 + O2 2SO3

Ni, hydrogenation of alkenes: e.g. H2C=CH2 + H2 H3C–CH3

MnO2, decomposition of hydrogen peroxide: 2H2O2 O2 + 2H2O

Co, as part of vitamin B12

Pt/Pd/Rh, used in catalytic converters

the periodic table 4 - pedagogics.ca · 2012-02-20 · 144 actinides lanthanides transition...

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