1 Opening items

1.1 Module introduction

One of the greatest triumphs of Western civilization has been the detailed understanding that scientists have obtained of everyday, large–scale phenomena in terms of the small–scale behaviour of atoms. Much of this understanding has been based on the theory of quantum mechanics developed in the mid-1920s, mainly by European physicists. However, the applications to atoms and molecules were based on ideas that had been formulated in the previous century by far–sighted chemists who understood more quickly than most physicists the reality of atoms and who appreciated that atomic ideas could be used to enable a profound understanding of the properties of the elements.

Most of this module is devoted to achieving an understanding of individual chemical elements and how they bond together to form compounds. This will lead us to consider some of the most important ideas in fundamental atomic science.

We begin by examining the first ionization energies of the elements and by demonstrating the information that these energies shed on the chemical properties of the elements. For example, the noble gases, which are very unreactive, have high values of first ionization energy, while the very reactive alkali metals have low values of first ionization energy. The relationship between the chemical reactivity of an atom and the arrangement of that atom’s electrons in various shells and subshells (its electronic structure) is the main theme of Section 2.

A periodicity in the chemical properties of the elements was noticed by the Russian chemist Dmitri Mendeleev, who arranged the elements into a periodic table to reflect this. The modern form of Mendeleev’s periodic table of the elements (Section 3) can be understood in terms of the distribution of electrons within the atom. The periodic table is arranged so that all elements with a similar outer shell of electrons (and hence similar chemical properties) appear in vertical columns, called groups, in which the principal quantum number of the outer shell of electrons increases down the column. The horizontal rows in the periodic table, the periods, typically contain several series of elements each of which is characterized by a common value of the principal quantum number. At both sides of the table are the columns of typical or main group elements. In between the main group elements are three series of transition elements (Subsection 3.2), a series of lanthanides (Subsection 3.3) and a series of actinides (Subsection 3.4).

Section 4 discusses the topic of chemical bonding. Simple theories of chemical bonding assume that the electronic configurations of the noble gases are especially stable, and that many elements tend to attain noble gas electronic configurations in chemical reactions. According to simple theories of bonding there are two ways in which atoms can attain this structure – electron transfer and electron sharing. This leads to the formation of ionic and covalent bonds, respectively, and these are discussed in Subsections 4.3, 4.4 and 4.5. Subsection 4.7 introduces a third kind of bonding, metallic bonding, in which each atom is presumed to share its bonding electron or electrons with every other atom in a crystal structure of positive ions.

Study comment Having read the introduction you may feel that you are already familiar with the material covered by this module and that you do not need to study it. If so, try the following Fast track questions. If not, proceed directly to the Subsection 1.3Ready to study? Subsection.

1.2 Fast track questions

Study comment Can you answer the following Fast track questions? If you answer the questions successfully you need only glance through the module before looking at the Subsection 5.1Module summary and the Subsection 5.2Achievements. If you are sure that you can meet each of these achievements, try the Subsection 5.3Exit test. If you have difficulty with only one or two of the questions you should follow the guidance given in the answers and read the relevant parts of the module. However, if you have difficulty with more than two of the Exit questions you are strongly advised to study the whole module.

Question F1

(a) What, in general, are the outer shell structures of noble gas atoms?

(b) Using your answer to part (a), write down the outer shell structure of the ground state of krypton, which is in the fourth period of the periodic table.

(c) Given that krypton has atomic number 36, write down the outer shell structure of the ground state of a bromine atom, which has atomic number 35. In which period and group of the periodic table would you expect bromine to lie?

Answer F1

(a) Noble gas atoms have an outer shell structure ns2np6, where n is the principal quantum number of the outer shell.

(b) 4s2 4p6. Since krypton is in the fourth period of the table, the principal quantum number of each electron in its outer shell is n = 4.

(c) The outer shell of the bromine atom will have the structure 4s2 4p5, because the atom contains one fewer electron than the krypton atom. Bromine is in group 7, the one next to the group of noble gas elements (group 0). It is, like krypton, in the fourth period because in both cases the principal quantum number of the outer shell is n = 4.

Question F2

What is the essential distinguishing feature of the electronic structures of each of the following series of elements: (a) transition elements, (b) lanthanides, (c) actinides?

(Assume that the elements are in their ground states.)

Answer F2

(a) In most transition elements, a d subshell is incomplete when the s subshell of the outer shell contains at least one electron (and usually two). To put it less formally, you could say that in the transition elements, inner d subshells begin to fill while an outer s subshell already contains at least one electron.

(b) In most of the lanthanides, the 4f subshells begin to fill when there are two electrons in the 6s subshell.

(c) In most of the actinides, the 5f subshells begin to fill when there are two electrons in the 7s subshell.

Question F3

What is the charge of the oxygen ion in magnesium oxide?

Answer F3

Oxygen has the electronic configuration 1s2 2s2 2p4. It can attain the noble gas electronic configuration (neon) by acquiring two electrons, that is, by forming an O2− ion. Oxygen in magnesium oxide therefore has a charge of −2 (where +1 is defined as the electric charge of the proton). Because chemical compounds are electrically neutral, the charges of the oxygen ion and the magnesium ion must balance each other. The formula of the oxide is therefore MgO, corresponding to Mg2+O2−.

Question F4

Sulphur (S) forms an oxide in which sulphur has a valency of six. What is the empirical formula of that oxide?

Answer F4

The empirical formula is SO3. The valency of oxygen is two, so as the valency of sulphur is six, the ratio of sulphur atoms to oxygen atoms is 1 : 3.

Question F5

Examine statements (a) to (c) below, which describe properties of carbon tetrachloride (CCl4), fluorine oxide (F2O), and rubidium oxide (Rb2O). These substances are either ionic or covalent. Where you think a substance is ionic, assign charges to each ion. Where you think a substance is covalent, write a Lewis structure for it.

(a) CCl4 is a liquid at room temperature and does not conduct electricity; it boils at 77 °C.

(b) F2O is a gas at room temperature.

(c) Rb2O is a solid with a high melting temperature. The structure of this oxide is such that each rubidium ‘unit’ is surrounded by four equidistant oxygen ‘units’, while each oxygen ‘unit’ is surrounded by eight equidistant rubidium ‘units’.

Answer F5

(a) and (b) CCl4 and F2O are covalent compounds. The low melting and boiling temperatures suggest discrete molecular compounds and, in the case of CCl4, this is corroborated by the low conductivity of the liquid. For CCl4 and F2O we can write the following Lewis structures:

(c) Rb2O is an ionic compound. This is apparent from its high melting temperature and from its structure: each oxygen is surrounded by eight rubidium ‘units’ while each rubidium ‘unit’ is surrounded by four oxygen ‘units’. Such structures are characteristic of ionic compounds.

In the case of Rb2O, the ions Rb+ and O2− have noble gas electronic configurations (krypton and neon, respectively). We can therefore write Rb2O as (Rb+)2O2−.

Study comment Having seen the Fast track questions you may feel that it would be wiser to follow the normal route through the module and to proceed directly to the following Ready to study? Subsection.

Alternatively, you may still be sufficiently comfortable with the material covered by the module to proceed directly to the Section 5Closing items.

1.3 Ready to study?

Study comment In order to study this module you will need to understand the following terms: atom, atomic number, chemical compound, electron, electric charge, electric current, element, ion, kelvin (the SI unit of temperature), molecule, nucleus, proton and voltage. You should also be familiar with the use of chemical symbols to represent the chemical elements, and their combinations in chemical formulae, such as H2O. (The symbols and atomic numbers of the elements are shown in Table 2). If you are uncertain about any of these terms, review them now by reference to the Glossary, which will also indicate where in FLAP they are developed. The following Ready to study questions will allow you to establish whether you need to review some of the topics before embarking on this module.

Question R1

The atomic number of argon is Z = 18. What information can you deduce about a neutral argon atom from this fact?

Answer R1

Every atom consists of a positively charged central nucleus surrounded by negatively charged electrons. An argon atom with Z = 18 has a nuclear charge that is 18 times greater than that on a hydrogen nucleus (i.e. 18e, where e is the charge on the proton). In order that it should be electrically neutral, an argon atom must have 18 electrons ‘orbiting’ its nucleus. Apart from revealing the charge on the nucleus and the number of electrons in the neutral atom, the atomic number provides other items of information regarding X–ray spectra, etc. but these may be seen as consequences of the charge on the nucleus.

Consult the relevant terms in the Glossary for further information.

Question R2

In terms of their electrical charges, what is the key difference between a lithium ion and a helium atom?

(The atomic numbers of lithium and helium are 3 and 2, respectively.)

Answer R2

A lithium ion is positively charged, whereas a helium atom is neutral: lithium has three protons in its nucleus, whereas helium has only two. The superscript on the symbol for the lithium ion (Li+) records the single excess positive charge.

Consult atomic number and ion in the Glossary for further information.

2 Chemical evidence for electron shells in atoms

In this section, we will consider patterns that are evident in the behaviour of chemical elements and how these patterns are related to the underlying atomic structures of the elements. We are thus following a standard scientific path by trying to understand behaviour on the large scale in terms of structure on a much smaller scale.

2.1 Periodicity in chemical reactivity

Figure 1 First ionization energies of elements up to radon.

Let us begin with atoms, which generally consist of a tiny central nucleus that is in some sense ‘orbited’ by electrons (only hydrogen atoms contain a lone electron). When all the electrons in an atom have their lowest possible energy, the atom as a whole has its lowest energy and is said to be in its ground state. This is a very convenient ‘reference state’ and it is safe to assume that whenever the structure or behaviour of an atom is described in this module, that atom will be in its ground state unless you are told otherwise.

The minimum energy required to completely remove an electron from an atom in its ground state is called the ionization energy, a quantity that can readily be measured in the laboratory. The energy required to completely remove the least tightly bound electron from the influence of the atom is called the element’s first ionization energy. This quantity is plotted against atomic number in Figure 1 for each element up to radon, Rn (Z = 86). Look carefully at this plot – it contains several important clues about atomic structure as we will now describe.

You can see from Figure 1 that the first ionization energies form a regular pattern: there are a number of peaks, each followed by a deep trough. At the peaks are the elements helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn). These elements, all gases at room temperature, are all inert, that is, they are very reluctant to combine with other elements. For this reason, they are collectively referred to as the noble gases (sometimes known as the inert gases).

The elements in the troughs that immediately follow the noble gases are lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). These have similar properties to each other, but are very different from the noble gases: they are all highly reactive metals and are called the alkali metals because they react violently (or even explosively!) with water to form substances called alkalies. (These alkalies are compounds of an alkali metal with oxygen and hydrogen; examples are LiOH, NaOH and KOH.)

Now consider the elements that occur just before the peaks in Figure 1 – fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At). These elements, called the halogens, i also have very similar chemical properties: each of them has a rather unpleasant odour and combines readily with many other elements. For example, each halogen combines with hydrogen to form an acid. (Narrowly defined, acids are compounds that include hydrogen and which form hydrogen ions when dissolved in water.) i

The pattern in first ionization energies plainly indicates a corresponding pattern in the chemical behaviour of the elements. Long before ionization energies were measured or even thought of – indeed, well before it was realized that the atoms generally consist of a nucleus and electrons – several scientists had noticed a pattern in the chemical properties of the elements. The most important and famous insight was provided by the great Russian chemist Dmitri Mendeleev (1834–1907), who first arranged the elements in order of increasing atomic mass to form a periodic table.

To see why the idea of such a table seemed plausible to him, consider first the elements with atomic numbers from 2 to 9 inclusive:

He Li Be B C N O F

Although Mendeleev did not know their atomic numbers or anything about ionization energies, he did know the chemical formulae of some of the compounds formed by these elements, notably those with oxygen (oxides) and hydrogen (hydrides). Consider first the oxides of these elements. If we write the formulae of the oxides as M2On, where M is any element from the list above, then from helium to boron (B), n takes the values 0, 1, 2 and 3 (the value zero for helium is appropriate because it does not form an oxide). The next two elements, carbon (C) and nitrogen (N), form more than one oxide. However, if you pick out the highest oxide, i.e. the oxide that contains the highest proportion of oxygen, you obtain n = 4 for carbon and n = 5 for nitrogen. Thus, the trend that runs from helium to boron continues to nitrogen. Beyond nitrogen, however, it peters out: F2O7, for example, does not exist.

In the case of hydrides, we can write the formula MHn. Again if we restrict ourselves to the highest hydrides, i.e. those that contain the greatest proportion of hydrogen, then right across the row, n follows the sequence 0, 1, 2, 3, 4, 3, 2, 1: the values of n rise in steps of one until they reach a maximum value of four, before falling off in steps of one unit. Finally, the formulae of the highest chlorides (compounds formed with chlorine) follow the same trend as those of the highest hydrides.

Suppose we now look for patterns in the next eight elements, which we will write underneath the ones we have just discussed:

He Li Be B C N O F

Ne Na Mg Al Si P S Cl

This time the oxide trend is perfect. When the formulae of the highest oxides are written as M2On, n takes the values 0, 1, 2, 3, 4, 5, 6, 7, from neon to chlorine. The trend in the formulae of the highest hydrides is exactly the same as in the first row, and so is that in the formulae of chlorides, except that at phosphorus (P), the higher chloride PCl5, is formed. In this type of arrangement, elements with similar chemical properties (helium and neon, lithium and sodium) fall in vertical columns called groups. The horizontal rows are called periodic_seriesperiods.

In Mendeleev’s periodic table the group numbers are meant to represent the number of oxygen atoms that two atoms of the element are combined with in the highest oxide. For example, elements in group 2 form highest oxides of general formula MO, with two oxygen atoms combined with every two atoms of M.

2.2 Reactivity determined by outer electron configuration

Study comment This subsection includes a highly condensed introduction to some of the essential features of the electronic configuration of atoms. A more thorough and less condensed treatment of this topic is given elsewhere in FLAP. See the Glossary for details.

We have now established the structure of the periodic table, which neatly summarizes regularities in the chemical behaviour of chemical elements. In our discussion, we have often referred to the reactivity of the elements – how readily they combine with others to form new substances – but we have not yet explained why certain elements react in similar ways. This is the task of this subsection: once again, we will be explaining large–scale phenomena (chemical reactivity) in terms of small–scale structure.

An atom generally consists of a nucleus surrounded by electrons, so initial contact between atoms occurs in the vicinity of the outer electrons. Not surprisingly, therefore, it is the outer electronic configuration that strongly influences the way one kind of atom bonds to another.

Electrons in atoms are arranged in shells, characterized by the principal quantum number n (n = 1 corresponds to the first shell, n = 2 to the second, and so on); each shell contains one or more subshells, characterized by the orbital angular momentum quantum number l. In the nth shell there are n possible subshells, each of which corresponds to an integer (i.e. whole number) value of l in the range 0 to n – 1. The subshells corresponding to l values of 0, 1, 2 and 3 are often denoted by the letters s, p, d and f, respectively: so an electron with n = 1 and l = 0 is called a 1s electron, one with n = 2 and l = 1 a 2p electron, and so on. The quantum numbers of an electron are said to determine its quantum state and this, in turn, determines its energy level. To a first approximation, in any given atom, all the electrons in a given subshell have the same energy.

✦ What are the quantum numbers of an electron in the 3d state?

n = 3, l = 2. The electron in the 3d state has the principal quantum number n = 3 and its orbital angular momentum quantum number is l = 2, in common with all d state electrons.

As we move from one element to another in order of increasing atomic number, the number of electrons increases. Fundamental principles of atomic physics, (particularly Pauli’s exclusion principle), imply that the maximum number of electrons that can simultaneously occupy a subshell characterized by an orbital angular momentum quantum number l is 2(2l + 1). Hence no s subshell can contain more than two electrons, no p subshell more than six, no d subshell more than ten and no f subshell more than 14. This limitation plays a major part in determining the distribution of electrons across the various shells and subshells – the electronic configuration of the atom.

In particular, the ground state electronic configurations, in which the electrons are distributed so as to minimize the atom’s energy, tend to become more complicated as the atomic number increases. If we examine the precise way in which the electronic configuration changes as the atomic number increases we find that this too shows a periodic variation.

To see this let us consider part of the sequence of elements that we looked at in Subsection 2.1:

Li Be B C N O F Ne

(You may have noticed that helium has disappeared; it will come back later!)

The (ground state) electronic configuration of the lithium atom has two electrons in the s subshell of the n = 1 shell (the maximum number allowed) and one electron in the s subshell of the n = 2 shell. We represent this structure by writing it as 1s2 2s1, the full size numbers indicate shells, the letters the subshells, and the superscripts show the number of electrons in each subshell. In beryllium there is a second 2s electron, so the electronic structure is 1s2 2s2. No s subshell can contain more than two electrons, so in the boron atom the fifth electron must be a 2p electron, and the electronic configuration of boron is 1s2 2s2 2p1. From boron to neon, successive electrons are added to the 2p level, until at neon the configuration is 1s2 2s2 2p6. So the electronic configurations of the elements from lithium to neon are as shown in the margin. i

At neon, the 2p level is full because it contains its maximum quota of six electrons. The next highest energy level is 3s so the configuration of the element sodium, which contains one more electron than neon, is 1s2 2s2 2p6 3s1. For the seven elements after sodium, the 3s and 3p levels are filled in the same way that the 2s and 2p levels are filled across the second period.

✦ What is the electronic configuration of the elements magnesium and aluminium, which contain respectively one and two more electrons than sodium?

✧ For magnesium, the configuration is 1s2 2s2 2p6 3s2. As you have just seen, the electronic configuration of sodium is 1s2 2s2 2p6 3s1. The additional electron will go into the 3s state. For aluminium, the structure is 1s2 2s2 2p6 3s2 3p1 because the next electron must go into the 3p state (it cannot go into the 3s state, which can only accommodate a maximum of two electrons, in common with all s states).

You can therefore group the two rows (periods) as follows:

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

As you can see, this is very similar to the way that Mendeleev arranged the elements, except that the noble gases neon and argon are now at the ends of the rows whereas in our earlier arrangement neon was at the beginning of a row. This explains the mysterious disappearance of helium from our sample of elements – helium is an noble gas, like neon and argon, and therefore in our present arrangement should appear on the right of the table, above neon:

              He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

As we will describe shortly, this arrangement makes more sense in terms of electronic configuration and is the one invariably adopted in modern versions of the periodic table.

✦ What common feature is shared by the electronic configurations of elements in the same vertical column?

✧ Their outer electronic configurations i are the same except that the shell number changes. For example, for both carbon and silicon, the occupied level of highest energy is a p subshell and in both cases it contains two electrons.

Now let us pause to summarize what we have covered so far. The electronic configuration of atoms is a periodic property because when we set out the elements in order of atomic number, elements whose atoms have similar outer electronic configurations, such as carbon and silicon, occur at regular intervals. By grouping such elements in columns, we are constructing a periodic table that represents the periodicity in electronic configuration.

The elements in the two rows that we have set out have outer electronic configurations of the type ns1, ns2 and ns2np1 to ns2np6. Such elements are called typical elements or main group elements.

It is now plain that the members of certain families of elements which we introduced earlier have similar outer electronic configurations and are included among the typical elements. Thus, the alkali metals, which include lithium and sodium, have outer electronic configurations of the type ns1. Again, the noble gases, with the exception of helium, have outer electronic configurations of the type ns2np6.

The electronic configuration of these elements is shown in Table 1. They are known as group 0 elements.

Table 1 Ground states of group 0 elements (noble gases). All group 0 elements
(except helium) have outer shell configuration of ns2np6.
Element Symbol Outer shell
helium He 1s2
neon Ne 1s2 2s2 2p6
argon Ar 1s2 2s2 2p6 3s2 3p6
krypton Kr 1s2 2s2 2p6 3s2 3p6 3d104s2 4p6
xenon Xe 1s2 2s2 2p6 3s2 3p6 3d104s2 4p6 4d105s2 5p6
radon Rn 1s2 2s2 2p6 3s2 3p6 3d104s2 4p6 4d104f145s25p6 5d106s2 6p6

Because these elements have the maximum possible number of electrons in their highest–energy quantum states (i.e. they each have a full outer subshell), they are very unreactive.

The number of electrons in the outer subshell increases across a period from left to right across the periodic table, and so the natural place for this particular group of noble gases is at the extreme right–hand side of the table.

3 The periodic table

The formulation of the periodic table was undoubtedly one of the most important events in the history of science. It introduced a powerful framework for all subsequent work in chemistry – a colossal achievement in itself – and it brought order to a bewildering diversity of data on the elements. In this section, we will look more closely at the periodic table and you will be able to see clearly its power as a means of organizing information about the chemical elements.

Figure 2 The periodic table of the elements.

Table 2a Ground state electronic configurations of the elements. Typical elements (i).
Z   1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s
1 H 1                                  
2 He 2                                  
3 Li 2 1                                
4 Be 2 2                                
5 B 2 2 1                              
6 C 2 2 2                              
7 N 2 2 3                              
8 O 2 2 4                              
9 F 2 2 5                              
10 Ne 2 2 6                              
Table 2b Ground state electronic configurations of the elements. Typical elements (ii).
Z   1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s
11 Na 2 2 6 1                            
12 Mg 2 2 6 2                            
13 Al 2 2 6 2 1                          
14 Si 2 2 6 2 2                          
15 P 2 2 6 2 3                          
16 S 2 2 6 2 4                          
17 Cl 2 2 6 2 5                          
18 Ar 2 2 6 2 6                          
19 K 2 2 6 2 6   1                      
20 Ca 2 2 6 2 6   2                      
Table 2c Ground state electronic configurations of the elements. Transition elements (i).
Z   1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s
21 Sc 2 2 6 2 6 1 2                      
22 Ti 2 2 6 2 6 2 2                      
23 V 2 2 6 2 6 3 2                      
24 Cr 2 2 6 2 6 5 1                      
25 Mn 2 2 6 2 6 5 2                      
26 Fe 2 2 6 2 6 6 2                      
27 Co 2 2 6 2 6 7 2                      
28 Ni 2 2 6 2 6 8 2                      
29 Cu 2 2 6 2 6 10 1                      
30 Zn 2 2 6 2 6 10 2                      
Table 2d Ground state electronic configurations of the elements. Typical elements (iii).
Z   1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s
31 Ga 2 2 6 2 6 10 2 1                    
32 Ge 2 2 6 2 6 10 2 2                    
33 As 2 2 6 2 6 10 2 3                    
34 Se 2 2 6 2 6 10 2 4                    
35 Br 2 2 6 2 6 10 2 5                    
36 Kr 2 2 6 2 6 10 2 6                    
37 Rb 2 2 6 2 6 10 2 6     1              
38 Sr 2 2 6 2 6 10 2 6     2              
Table 2e Ground state electronic configurations of the elements. Transition elements (ii).
Z   1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s
39 Y 2 2 6 2 6 10 2 6 1   2              
40 Zr 2 2 6 2 6 10 2 6 2   2              
41 Nb 2 2 6 2 6 10 2 6 4   1              
42 Mo 2 2 6 2 6 10 2 6 5   1              
43 Tc 2 2 6 2 6 10 2 6 6   1              
44 Ru 2 2 6 2 6 10 2 6 7   1              
45 Rh 2 2 6 2 6 10 2 6 8   1              
46 Pd 2 2 6 2 6 10 2 6 10                  
47 Ag 2 2 6 2 6 10 2 6 10   1              
48 Cd 2 2 6 2 6 10 2 6 10   2              
Table 2f Ground state electronic configurations of the elements. Typical elements (iv).
Z   1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s
49 In 2 2 6 2 6 10 2 6 10   2 1            
50 Sn 2 2 6 2 6 10 2 6 10   2 2            
51 Sb 2 2 6 2 6 10 2 6 10   2 3            
52 Te 2 2 6 2 6 10 2 6 10   2 4            
53 I 2 2 6 2 6 10 2 6 10   2 5            
54 Xe 2 2 6 2 6 10 2 6 10   2 6            
55 Cs 2 2 6 2 6 10 2 6 10   2 6     1      
56 Ba 2 2 6 2 6 10 2 6 10   2 6     2      
Table 2g Ground state electronic configurations of the elements. The Lanthanides.
Z   1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s
57 La 2 2 6 2 6 10 2 6 10   2 6 1   2      
58 Ce 2 2 6 2 6 10 2 6 10 2 2 6     2      
59 Pr 2 2 6 2 6 10 2 6 10 3 2 6     2      
60 Nd 2 2 6 2 6 10 2 6 10 4 2 6     2      
61 Pm 2 2 6 2 6 10 2 6 10 5 2 6     2      
62 Sm 2 2 6 2 6 10 2 6 10 6 2 6     2      
63 Eu 2 2 6 2 6 10 2 6 10 7 2 6     2      
64 Gd 2 2 6 2 6 10 2 6 10 7 2 6 1   2      
65 Tb 2 2 6 2 6 10 2 6 10 9 2 6     2      
66 Dy 2 2 6 2 6 10 2 6 10 10 2 6     2      
66 Ho 2 2 6 2 6 10 2 6 10 11 2 6     2      
68 Er 2 2 6 2 6 10 2 6 10 12 2 6     2      
69 Tm 2 2 6 2 6 10 2 6 10 13 2 6     2      
70 Yb 2 2 6 2 6 10 2 6 10 14 2 6     2      
Table 2h Ground state electronic configurations of the elements. Transition elements (iii).
Z   1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s
71 La 2 2 6 2 6 10 2 6 10 14 2 6 1   2      
72 Ce 2 2 6 2 6 10 2 6 10 14 2 6 2   2      
73 Pr 2 2 6 2 6 10 2 6 10 14 2 6 3   2      
74 Nd 2 2 6 2 6 10 2 6 10 14 2 6 4   2      
75 Pm 2 2 6 2 6 10 2 6 10 14 2 6 5   2      
76 Sm 2 2 6 2 6 10 2 6 10 14 2 6 6   2      
77 Eu 2 2 6 2 6 10 2 6 10 14 2 6 9          
78 Gd 2 2 6 2 6 10 2 6 10 14 2 6 9   1      
79 Tb 2 2 6 2 6 10 2 6 10 14 2 6 10   1      
80 Dy 2 2 6 2 6 10 2 6 10 14 2 6 10   2      
Table 2i Ground state electronic configurations of the elements. Typical elements (v).
Z   1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s
81 Tl 2 2 6 2 6 10 2 6 10 14 2 6 10   2 1    
82 Pb 2 2 6 2 6 10 2 6 10 14 2 6 10   2 2    
83 Bi 2 2 6 2 6 10 2 6 10 14 2 6 10   2 3    
84 Po 2 2 6 2 6 10 2 6 10 14 2 6 10   2 4    
85 At 2 2 6 2 6 10 2 6 10 14 2 6 10   2 5    
86 Rn 2 2 6 2 6 10 2 6 10 14 2 6 10   2 6    
87 Fr 2 2 6 2 6 10 2 6 10 14 2 6 10   2 6   1
88 Ra 2 2 6 2 6 10 2 6 10 14 2 6 10   2 6   2
Table 2j Ground state electronic configurations of the elements. The Actinides.
Z   1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s
89 Ac 2 2 6 2 6 10 2 6 10 14 2 6 10   2 6 1 2
90 Th 2 2 6 2 6 10 2 6 10 14 2 6 10   2 6 2 2
91 Pa 2 2 6 2 6 10 2 6 10 14 2 6 10 2 2 6 1 2
92 U 2 2 6 2 6 10 2 6 10 14 2 6 10 3 2 6 1 2
93 Np 2 2 6 2 6 10 2 6 10 14 2 6 10 4 2 6 1 2
94 Pu 2 2 6 2 6 10 2 6 10 14 2 6 10 6 2 6   2
95 Am 2 2 6 2 6 10 2 6 10 14 2 6 10 7 2 6   2
96 Cm 2 2 6 2 6 10 2 6 10 14 2 6 10 7 2 6 1 2
97 Bk 2 2 6 2 6 10 2 6 10 14 2 6 10 8 2 6 1 2
98 Cf 2 2 6 2 6 10 2 6 10 14 2 6 10 10 2 6   2
99 Es 2 2 6 2 6 10 2 6 10 14 2 6 10 11 2 6   2
100 Fm 2 2 6 2 6 10 2 6 10 14 2 6 10 12 2 6   2
101 Md 2 2 6 2 6 10 2 6 10 14 2 6 10 13 2 6   2
102 Nu 2 2 6 2 6 10 2 6 10 14 2 6 10 14 2 6   2
Table 2k Ground state electronic configurations of the elements. Transition elements (iv).
Z   1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s
103 Lr 2 2 6 2 6 10 2 6 10 14 2 6 10 14 2 6 1 2
104 Rf 2 2 6 2 6 10 2 6 10 14 2 6 10 14 2 6 2 2
105 Ha 2 2 6 2 6 10 2 6 10 14 2 6 10 14 2 6 3 2

3.1 Noble gases, halogens and alkali metals

Figure 2 shows the modern version of the periodic table. You will probably find it helpful to look at this figure in conjunction with Table 2, which shows the electronic configurations of the elements (up to element 105). i

The periodic table is arranged so that all elements with a similar outer shell electronic configuration (and hence similar chemical properties) appear in vertical columns, groups, in which the principal quantum number of the outer shell increases as you step down the column. Thus the horizontal rows (periods) in the periodic table contain all the elements with the same principal quantum number in the outer shell. The subshells fill up as you move from left to right across the rows.

We can now discuss the properties of the elements in the light of their position in the table. The noble gases, the group 0 elements, appear in a column down the right hand side of the table. In Subsection 2.1 we mentioned two of the groups in the periodic table which contain particularly reactive elements. The group 7 elements, known as halogens, all have outer shell electronic configuration ns2np5, as you would expect since they appear one column to the left of the noble gases. Thus, all group 7 elements have one vacant space in their outer p subshell. This space is easily filled by a spare electron resulting in a negatively charged ion which is then very reactive on account of its electrostatic properties. The group 1 elements on the other hand (the alkali metals) all have an outer shell structure ns1. This outer electron is weakly bound, which is manifest in the low ionization energies of these elements, and hence easily removed to leave a positively charged ion. Very often, of course, positively charged alkali metal ions combine with negatively charged halogen ions to form salts such as sodium chloride (NaCl), the chemical name for table salt.

3.2 The transition elements

The two elements that follow argon in the table_2bperiodic table are potassium and calcium. These are typical elements, because their outer electronic configurations are 4s1 and 4s2, respectively. Potassium can be grouped under sodium, and calcium under magnesium. The next element is scandium Sc, which is certainly not a typical element: when the 4s level is full, the 3d level begins filling rather than the 4p. As table_2cTables 2c, table_2e2e, table_2h2h and table_2k2k show, this pattern of the d subshell in the nth shell filling after the s subshell in the n + 1th shell is repeated in each of the later periods in the periodic table.

Figure 3 gives the subshell energy level diagram and Figure 4 displays the order of ascending subshells.

Figure 4 Order of ascending energy for the subshells.

Figure 3 Subshell diagram.

If we write out the subshells in the order of increasing energy, then the electronic configuration of scandium is 1s2 2s2 2p6 3s2 3p6 4s2 3d1. The 4p level only begins filling at gallium Ga (element 31), when the 3d level has been filled. Typical elements thus reappear in the table at gallium (Table 2d).

The elements between calcium and gallium (Table 2c) span the region in which the 3d shell is filling up. There are ten of them, because the 3d level can accommodate up to ten electrons. Because their configurations are of this type, they cannot be grouped under typical elements, so the block of typical elements is usually separated between calcium and gallium, and the ten elements are inserted there.

The elements scandium Sc (Z = 21) to zinc Zn (Z = 30) are called transition elements because they form a bridge between two blocks of typical elements. We define transition elements as the elements that span a region in which a d shell is filling.

✦ In the scandium to zinc series, what is the most common outer electronic configuration of the transition elements?

✧ It is 4s2 3dx, where x runs from 1 to 10 inclusive. For example, the outer electronic configuration of titanium (Ti) is 4s2 3d2.

Two of the ten transition elements, chromium (Cr) and copper (Cu), do not have outer electronic configurations of the type 4s2 3dx. Their outer electronic configurations are, respectively, 4s13d5 and 4s13d10. We will not go into the detailed causes of these exceptions, but they reflect the fact that, in the scandium to zinc region, the 3d and 4s energy levels are very close in energy.

The periodic table in Figure 2 clearly sets the transition elements apart from the typical elements, and there is strong experimental support for this – the transition elements have properties different from those of the typical elements. Among the characteristics of transition elements are the following:

(a)

They are all metals.

(b)

Their compounds are often strongly coloured (for example, many copper compounds are a characteristic blue–green and many manganese compounds are pink).

(c)

They often form more than one compound with other elements. In particular, they often form two or more halides with the formula type MXn (where M represents any metal and X represents a particular halogen element) and two or more oxides with the formula type M2On. In such cases the values of n usually differ by one. Thus iron forms a dichloride FeCl2, and a trichloride FeCl3, which are pale green and yellow-brown, respectively. Chromium forms the following fluorides: CrF2 (green), CrF3 (green), CrF4 (green), CrF5 (red) and CrF6 (yellow), and copper yields the oxides Cu2O and CuO. By contrast, typical elements form two or more chlorides or oxides less often, and when they do the values of n usually differ by two. Thus lead forms just two chlorides, PbCl2 and PbCl4.

(d)

Magnetism is often associated with transition elements. For example, the strongly magnetic metals (Fe, Co, Ni) all lie in the first transition region.

After the series of transition elements has been completed, as we have said, the 4p level begins filling at gallium (Z = 31). When the 4p level becomes full at the noble gas krypton (Z = 36), the fourth period ends.

In terms of the electronic configurations of its elements, the fifth period is a straightforward repeat of the fourth. From Figures 3 and 4 you can see that the 5s level is filled first (rubidium, Rb, and strontium, Sr), after which the filling of the 4d level generates a second series of transition elements, from yttrium (Y) to cadmium (Cd).

Typical elements reappear with the filling of the 5p level at indium (In), and the period ends when the 5p level is full at the noble gas xenon.

3.3 The lanthanides

The sixth period begins with the elements caesium (Cs) and barium (Ba). If we represent the electronic configuration of the noble gas, xenon, by the symbol [Xe], then the electronic configurations of caesium and barium are [Xe]6s1 and [Xe]6s2, respectively. The atomic number of barium is 56.

Question T1

Use Figure 3 to predict the electronic configurations of the following elements (their atomic numbers are given in brackets): lanthanum (57), praseodymium (59), terbium (65), ytterbium (70), lutetium (71).

Answer T1

From Figure 3 you can see that after 6s, the level of next highest energy is 4f, followed by 5d. The 4f level can accommodate up to fourteen electrons (in common with all f states). The procedure we have been following therefore suggests that lanthanum should have the structure [Xe]6s2 4f1 because this atom contains one more electron than a barium atom and this should go into the next highest energy state, the 4f state.

Similarly, the electronic configuration of praseodymium should be [Xe]6s2 4f3 because the additional three electrons should all be accommodated in the 4f state. Using identical reasoning, the structures of the remaining elements should be: terbium, [Xe]6s2 4f9; ytterbium, [Xe]6s2 4f14 and lutetium, [Xe]6s2 4f14 5d1.

Comment With only one exception these predictions are correct; the exception is lanthanum, which in fact has the configuration [Xe]6s2 5d1.

The answer to Question T1 suggests that, after barium, a new series of elements appears in which the 4f shell is being filled. There will be 14 elements in this new series because f shells can accommodate 14 electrons. The 4f shell becomes full at ytterbium, which has the configuration [Xe]6s2 4f14.

We will therefore define the new series as the fourteen elements from lanthanum to ytterbium, inclusive. The fourteen elements that follow barium are called lanthanides and their characteristic electronic configuration is [Xe]6s2 4fx.

In general then, we can say that the lanthanides are typified by the outer electronic configuration 6s2 4f1x, where x runs from 1 to 14. i

Because they cannot be grouped beneath elements that we have considered up until now, we must pull the periodic table apart again, and insert them after barium (Figure 2). After the 4f level has been filled at ytterbium (Yt), lutetium (Lu), which has the configuration [Xe]6s2 4f14 5d1, begins a new series of transition elements.

The lanthanides behave very similarly in many chemical reactions, so similarly that only since 1945 has it become possible to separate them completely from each other and obtain pure samples of each.

For nearly all of them, the highest oxide has formula M2O3, and for all fourteen of them, the highest chloride is MCl3. i

As we have already stated, after the 4f level has been filled at ytterbium, lutetium begins a new series of transition elements. The new transition series is completed with element number 80, mercury (Hg), and more typical elements then appear as the 6p level begins to fill. When the 6p level is filled at the noble gas radon, the sixth period is complete.

The lanthanides and transition elements play a key role in high–temperature superconductivity, one of the most important areas of research to have begun in the 1980s. Superconductors are extremely unusual in that, below their characteristic critical temperatures, they can carry electricity with zero resistance. The phenomenon was discovered in 1911 by the Dutch physicist Heike Kammerlingh Onnes (1853–1926) who demonstrated that lead and mercury were superconductors when they are cooled to the temperature of liquid helium (4.2 K).

For almost 75 years, there seemed little hope of finding materials that were superconductors at temperatures much higher than about 23 K. However, in 1986 K. Alex Müller (1927–) and Georg Bednorz (1950–) discovered at the IBM research laboratories in Zurich that a certain form of barium lanthanum copper oxide was a superconductor up to about 35 K. Soon afterwards, Ching–Wu Chu and his team at the University of Houston discovered that a related class of ceramic materials, a particular form of yttrium barium copper oxide exhibited superconductivity up to an astonishing 84 K, well above the temperature of liquid nitrogen. Because liquid nitrogen is comparatively inexpensive (it is cheaper than beer!), these discoveries hold the promise of making superconductors available for commercial use.

Bednorz and Müller, the midwives of the new field of high–temperature superconductivity, were jointly awarded the 1987 Nobel Prize for Physics.

3.4 The actinides

The seventh period begins as usual with an alkali metal, francium (Fr), and a group 2 metal, radium (Ra), as the 7s level is filled. The elements beyond this point are classified with the lanthanide elements. This similarity of the elements beyond radium to the lanthanides became obvious only during the wartime programme for the manufacture of atomic bombs in the United States, when teams of scientists under the direction of Glen Seaborg (1912–) succeeded in making new elements with atomic numbers greater than that of uranium (U).

Among these elements is element number 94, plutonium (Pu), which can be made by neutron bombardment of uranium in a nuclear reactor.

At a certain point beyond uranium, marked by atomic numbers 95 and 96 (americium, Am, and curium, Cm), it was found that the elements began to behave like lanthanides. Thus americium and the elements immediately following it all form a highest chloride of formula MCl3.

This analogy with the lanthanides prompted Seaborg to propose in 1945 that actinium, and the thirteen elements following it, spanned a region in which the 5f level is filling. (Notice that this is consistent with Figure 4.) These fourteen elements are therefore grouped beneath the lanthanides in Figure 2; they are called the actinides.

Further confirmation of this idea has come with the preparation of elements 104–107. These were made by Soviet and American scientists between 1976 and 1986. These elements behave like transition elements rather than like actinides or lanthanides. This suggests that the 5f level is full at nobelium (No), element 102, and that stepwise filling of the 6d level begins at element 103, lawrencium (Lr) (Figure 4).

Question T2

(a) Use Figures 3 and 4 to work out the electronic configurations of the elements that have atomic numbers 16 and 28, respectively.

(b) Are these typical elements, transition elements, lanthanides or actinides?

(c) In what periods do they lie?

(d) Refer to the periodic table in Figure 2. In what group of the table does element 16 fall? What are the names of the two elements?

Answer T2

(a) The elements with atomic numbers 16 and 28 have electronic configurations 1s2 2s2 2p6 3s2 3p4 and 1s2 2s2 2p6 3s2 3p6 4s2 3d8, respectively. These configurations are determined by using Figure 4, which gives the order in which the subshells are filled. In order to find the configuration of the elements, it is simply necessary to bear in mind that s shells can accommodate up to two electrons, whereas p shells can accommodate up to six and d shells accommodate up to ten (this maximum occupancy is given at the top of Figure 3).

(b) Element 16 has an outer electronic configuration s2 px so it is a typical element; by contrast, element 28 has an outer electronic configuration s2 dx, which means that it is a transition element.

(c) The last subshells to be filled in element 16 are 3s and 3p, so Figure 3 tells you that the element is in period 3. In the same way, the last subshells to be filled in element 28 are 4s and 3d, so it is in period 4.

(d) Element 16 falls in group 6 of the periodic table: it is sulphur. Element 28 is nickel.

4 Chemical bonding

Our discussion so far has mainly concerned individual chemical elements and their properties. Now it is time to consider how they combine, or bond, to form other substances.

4.1 Valency

When two elements react to form a chemical compound, experiment shows that their atoms tend to combine in a particular ratio. For example, sodium and chlorine react to yield sodium chloride NaCl, not NaCl2 or NaCl3. Likewise, the product of the reaction of hydrogen and oxygen is H2O, not HO or HO2. Why is there such a strong preference for certain combinations such as NaCl and H2O?

Theories of chemical bonding supply answers to questions like this, and we will discuss them shortly. First, however, we will focus on the relationship between the numbers of atoms that combine in different compounds. An examination of this relationship leads to a property called valency, which is useful in predicting the formulae of certain compounds.

Consider the following three formulae:

HCl, NH3, NCl3

Notice that in HCl the elements hydrogen and chlorine combine with each other in the atomic ratio 3 : 1. Each of these two elements combines with nitrogen in the same ratio of 3 : 1 in NH3 and NCl3. There appears to be a relationship between the three formulae. This kind of observation for sets of compounds led chemists to suggest that the formula of a compound is determined by some inherent property of the constituent elements. This property, valency, determines the relative numbers of atoms of elements in compounds.

If each element had only a single value of its valency, the formulae of compounds could be readily predicted from a knowledge of those valencies. Unfortunately, many elements have several values of valency, which reduces the predictive power of the concept. Nevertheless, provided we steer clear of complicated cases, and show due caution, we can use the concept of valency to predict the formulae of many compounds quite straightforwardly.

The rules for determining the valency of an element are quite simple, particularly for binary compounds (those that contain only two elements), and we will restrict the discussion to these compounds. Because valency is linked to chemical formulae, it is concerned with the relative numbers of atoms in compounds. For this reason, some arbitrary reference element is chosen as the basis of values of valency. This element is hydrogen, and by definition its valency is taken to be one.

If an element forms a binary compound with hydrogen, its valency is equal to the number of atoms of hydrogen that combine with one atom of that element.

For example, in hydrogen chloride (HCl) one atom of chlorine combines with one atom of hydrogen, so the valency of chlorine is one.

✦ Nitrogen can combine with hydrogen to form the chemical ammonia, which has the chemical formula NH3. On the basis of this formula, what do you deduce to be the valency of nitrogen?

✧ In NH3, one atom of nitrogen combines with three atoms of hydrogen, so in this case the valency of nitrogen is three.

Knowing that nitrogen and chlorine have valencies of 3 and 1, respectively, (at least in the simple cases considered above), we can predict that they will jointly form the binary compound NCl3, which is indeed the case. The existence of NCl3 demonstrates that it is not necessary to rely on compounds of hydrogen in order to determine the valency of an element. For example, if the valency of chlorine in HCl is the same as that of hydrogen, i.e. one, the valency of nitrogen is equal to the number of atoms of chlorine that combine with one atom of nitrogen. In NCl3, therefore, the valency of nitrogen is three, the same value as we deduced from the existence of ammonia NH3.

To take another example, hydrogen and oxygen combine as water, which has the formula H2O.

Question T3

What is the valency of oxygen in H2O?

Answer T3

One atom of oxygen combines with two atoms of hydrogen, so the valency of oxygen is two.

Table 3 Valencies of ten common elements.
Element Valency
hydrogen (H) 1
carbon (C) 4
nitrogen (N) 3
oxygen (O) 2
fluorine (F) 1
sodium (Na) 1
magnesium (Mg) 2
chlorine (Cl) 1
potassium (K) 1
calcium (Ca) 2

As the examples above illustrate, the idea that elements have a fixed valency, is useful in predicting chemical formulae even though it must be treated with caution.

In general, if element A has valency x, and element B has valency y, then if A and B combine to form a compound, y atoms of A will combine with x atoms of B.

Of course, there are limitations to the generality of these predictions and to the information that valency provides. Valency alone gives no indication of the nature of the interaction between atoms, which is the subject of the following subsections. While you are reading them, you may find it helpful to refer to Table 3, which gives the valencies of ten common elements.

Question T4

In forming compounds, silicon (Si) generally exhibits a valency of four.

(a) What is the formula of the simplest substance that silicon can form with hydrogen?

(b) What is the formula of the simplest substance that silicon can form with oxygen (O), which has a valency of two?

Answer T4

(a) SiH4. If the valency of silicon is four, one atom of silicon can combine with four atoms of hydrogen.

(b) SiO2. One atom of silicon can combine with two atoms of oxygen.

4.2 A simple theory of chemical bonding

To form a compound, the different atoms have to bond together. An atom consists of a nucleus surrounded by electrons, so initial contact between atoms involves the outer electrons. Not surprisingly, therefore, the outer electronic configuration strongly influences the way one kind of atom bonds to another.

The simplest theories of chemical bonding are based on the observation that the elements of group 0, the noble gases, are extraordinarily unreactive. These elements are helium, neon, argon, krypton, xenon and radon. They are all monatomic gases: that is, they exist as single helium atoms or single neon atoms, not as He2 or Ne2 molecules. They do not react with the common laboratory chemicals. Indeed, they are all unaffected by molten sodium, which is usually extremely reactive, and helium, neon and argon do not even react with hot fluorine, a substance that attacks nearly every other element. Since 1962 some compounds of xenon, krypton and radon have been prepared, for example XeF4, XeO3 and KrF2, but the overall impression of a chemically very unreactive group is correct.

We describe this chemical inertness of the group 0 elements by saying that the noble gases are stable with respect to chemical reactions. Now when atoms react, their outer electronic configurations are disturbed. It follows that the electronic configurations of the noble gases must be especially stable to disturbance.

Figure 1 First ionization energies of elements up to radon.

As we noted at the beginning of this module, the ionization energies of the noble gases are unusually large (Figure 1), so removal of an electron from a noble gas atom is especially difficult. This suggests that the electronic configuration of a noble gas is particularly stable with respect to the loss of an electron.

Simple theories of chemical bonding assume that the electronic configurations of the noble gases are especially stable, and that many elements tend to attain noble gas electronic configurations in their chemical reactions.

According to this assumption, the noble gases are chemically inert because they already have the noble gas electronic configurations. How can one element attain the electronic configuration of another? According to the simple theories of bonding there are two ways: electron transfer and electron sharing.

4.3 Ionic bonding

The alkali metals, the group 1 elements, all have the electronic configuration of a noble gas plus one extra electron. If this electron is removed in a chemical reaction, a positively charged ion (called a cation) will be formed and this ion will have the electronic configuration of a noble gas atom. Thus, for example, lithium has the electronic configuration 1s22s1; by losing its 2s electron, it forms the lithium ion Li+ and attains the noble gas configuration of helium (1s2). This reaction can be written as:

Li → Li+ + e

Question T5

What is the electronic configuration of the sodium ion Na+? To what noble gas does this configuration correspond?

Answer T5

The configuration of Na+ is 1s2 2s2 2p6. This corresponds to the electronic configuration of neon.

Consider now the group 7 elements, the halogens. They all have the electronic configuration of a noble gas less one electron. By addition of one electron, for example in a chemical reaction, a negatively charged ion (called an anion) will be formed, which will have the electronic configuration of a noble gas atom. Thus, for example, fluorine (1s22s22p5) will form a fluoride ion F, with the electronic configuration (1s22s22p6) of neon.

Question T6

What is the electronic configuration of the chloride ion, Cl? To what noble gas does this configuration correspond? What does the superscript minus stand for?

Answer T6

The electronic configuration of Cl is 1s2 2s2 2p6 3s2 3p6. This corresponds to the configuration of the noble gas argon. The chloride ion now has 18 electrons (each with a single negative charge), but there are still only 17 positively charged protons in the nucleus. Consequently, the chlorine now has one excess negative charge, as the minus sign of Cl− illustrates.

Thus you can envisage that when the metal lithium and the gas fluorine react to form lithium fluoride (LiF), a colourless compound that resembles common salt, an electron is transferred from each lithium atom to each fluorine atom, with the formation of the following ions:

2Li + F2 → 2Li+ + 2F

The electrostatic attraction between the positive lithium ions and negative fluoride ions in lithium fluoride gives rise to ionic bonding.

We now know that lithium ions have a single positive charge, and fluoride ions have a single negative charge because these charges give each ion a noble gas electronic configuration. At the same time, this explains why lithium fluoride has the formula LiF: the compound must be neutral and this can only be achieved by combining equal numbers of the oppositely charged ions.

Question T7

Write down the electronic configuration of magnesium. How could magnesium most easily attain the electronic configuration of a noble gas? What therefore is the symbol for the magnesium ion? What is the formula for magnesium fluoride?

Answer T7

The electronic configuration of magnesium is 1s2 2s2 2p6 3s2. Loss of the two 3s2 electrons would produce a Mg++ ion (written as Mg2+) with the electronic configuration of neon, and two excess positive charges. So, in order to be neutral, magnesium fluoride must be MgF2, that is, Mg2+(F)2. This second formula records the fact that for every magnesium ion in the compound, there are two fluoride ions.

The alkali metals (group 1) form compounds containing singly charged positive ions, whereas magnesium in group 2 forms compounds containing the ion Mg2+, and the halogens (group 7) form compounds containing singly charged negative ions (generally called halide ions). All these ions have noble gas electronic configurations. Their charges are such that the alkali metals must form halides of the type M+X, and group 2 metals like magnesium, halides of the type M2+(X)2.

Ionic substances have a structure in which the ions tend to surround themselves with ions of opposite charge, and in which discrete molecules cannot be legitimately identified. Because there are strong electrostatic forces between the ions, ionic substances are generally solids at room temperature and they have comparatively high melting and boiling temperatures.

Another characteristic property of ionic substances is that when they melt they conduct electricity, because of the presence of charged ions. Also, if these substances dissolve significantly in water, they usually form solutions that readily conduct electricity.

4.4 Covalent bonding

Many substances, for example, oxygen (O2), carbon tetrachloride (CCl4) and ammonia (NH3), have properties that cannot be understood in terms of ions. Instead, they have a structure in which discrete molecules, with the same chemical formula as that of the substance, are separated from each other by what are relatively large distances on a molecular scale. The form of bonding that accounts for the existence of these molecules is known as covalent bonding, and the substances composed of such molecules are said to be molecular covalent substances.

At room temperature, these substances exist as solids, liquids and gases, and most of them have relatively low melting and boiling temperatures. They do not conduct electricity simply because they do not consist of ions.

Other examples of this type of substance are chlorine (Cl2), hydrogen (H2), nitrogen (N2), water (H2O), nitrogen trichloride (NCl3), methane (CH4) and carbon dioxide (CO2). Each of them comprise discrete molecules with the formula indicated. All of these formulae can be understood easily by assigning a fixed valency to the atoms of nitrogen, oxygen, hydrogen, chlorine and carbon: one to hydrogen and chlorine, two to oxygen, three to nitrogen and four to carbon.

The molecules H2, Cl2, NH3 and NCl3 can be represented as shown in structures I to IV:

In these molecular diagrams, the constituent atoms are linked by bonds that are represented by lines. The number of bonds issuing from each atom is equal to its valency, which determines the ways in which the atom can link with others. For example, because three bonds spring out of each nitrogen atom, whenever the nitrogen bonds with other atoms all three of its bonds must link with those of other atoms. One way in which this can be easily achieved is by bonding with three hydrogen atoms to form a molecule of ammonia, NH3.

Question T8

Represent molecules of the compounds hydrogen chloride (HCl), methane (CH4), carbon tetrachloride (CCl4) and water (H2O) using a diagram that makes clear the valencies of each constituent atom.

Answer T8

The four molecules are as given below:

The valencies are one for hydrogen and chlorine, two for oxygen and four for carbon. In each case, the number of lines issuing from an atom is equal to its valency.

Comment You may not have predicted the correct shape for the water molecule. The reasons for this shape are too complex to be discussed in this module.

The atoms in the molecular diagrams or structures that you have met so far have been linked by single valency lines or single bonds, but this is by no means always the case. Consider, for example, the oxygen (O2) and carbon dioxide (CO2) molecules. The molecules are drawn as in structures V and VI:

in order to ensure that oxygen has a valency of two and that carbon has a valency of four. In these examples, atoms are linked by two valency lines: they are said to be held together by double bonds.

✦ Draw a diagram of the N2 molecule in which nitrogen has its common valency of three.

✧ Here the two atoms are linked by a triple bond:

You now know how to draw diagrams of the molecules in molecular covalent compounds in which each atom forms a number of bonds equal to its assigned valency. What is the nature of the bonds which hold the atoms together, and why do the carbon, chlorine, hydrogen, nitrogen and oxygen atoms have their characteristic valencies?

The properties of chlorine, Cl2, are nothing like those of an ionic substance such as NaCl. However, even if we ignore this and try to give the Cl2 molecule an ionic form, we cannot do so consistently. For instance, suppose we say that one of the chlorine ‘units’ in the Cl2 molecule is a Cl ion with the electronic configuration of argon.

✦ What will be the electric charge and electronic configuration of the other chlorine ‘unit’?

✧ To maintain neutrality, it must be Cl+. This ion has the configuration of the sulphur atom.

In 1916, a solution to this sort of problem was proposed by Gilbert Lewis (1875–1946), the American chemist also famous for naming the photon (the quantum of electromagnetic radiation) in 1923. Lewis suggested that bonds could be formed between atoms by sharing pairs of electrons and that each shared pair constituted a covalent bond. He pointed out that this provided a means by which atoms in many compounds could attain noble gas structures. Thus, we can represent two chlorine atoms by the symbols given in structure VIII:

where the dots on one atom and the crosses on the other represent the seven outer electrons (3s23p5). We then write the chlorine molecule Cl2 as in IX:

This type of diagram is now called a Lewis structure. By convention, the outer electrons around the atoms are grouped in pairs. The pair of electrons that falls between the two chlorines is counted towards the electronic configurations of both atoms. The pair of electrons thus shared between the atoms constitutes a chemical bond.

✦ What electronic configuration is thereby achieved by the two chlorines in the Cl2 molecule?

✧ Each chlorine gains one electron. Both atoms thus attain the electronic configuration of argon, that is, they both have eight electrons in their outer shell (3s2 3p6 ).

Because one pair of electrons is involved in the Lewis structure, the bond is a single bond. The shared electron pair is therefore an electronic interpretation of the single bond line drawn between the two chlorine atoms in structure II.

As further examples, consider the molecules NCl3 and H2O. Nitrogen has five outer electrons (2s22p3), hydrogen has one (1s1), and oxygen has six (2s22p4). Suppose that we represent the outer electrons on the nitrogen and oxygen atoms by crosses, and those on the hydrogen and chlorine atoms by dots. Then each atom attains a noble gas electronic configuration in the Lewis structures are given in X and XI:

 i

✦ What noble gas electronic configurations are attained by nitrogen, oxygen and hydrogen in these structures?

✧ In NCl3, the nitrogen gains one electron from each of the three electron pairs that it shares with chlorine. Its eight outer electrons then correspond to the electronic configuration of neon (2s2 2p6). Here is the explanation of why the valency of nitrogen is three: this is the number of shared electron pairs needed to give nitrogen this noble gas electronic configuration.

The neon electronic configuration is also attained by oxygen in the water molecule: each hydrogen gains one electron from the single electron–pair bond that it forms with oxygen and attains the electronic configuration of helium (1s2).

Question T9

Draw a Lewis structure for the methane molecule, CH4.

Answer T9

Carbon has four outer electrons (2s2 2p2). It can attain the electronic configuration of neon (2s22p2) by gaining another four electrons from four shared electron pairs:

In this structure, each hydrogen again attains the helium electronic configuration.

So far, any two atoms in our Lewis structures have been held together by just one electron–pair bond. Sometimes, however, two or three pairs may be involved.

✦ Draw the Lewis structure corresponding to structure VI for CO2:

✧ Each atom can attain the noble gas electronic configuration of eight outer electrons if each oxygen is tied to the central carbon by not one but two electron–pair bonds. Therefore the Lewis structure will be as in XII.

The presence of two electron pairs between each oxygen ‘unit’ and the central carbon ‘unit’ allow us to interpret each carbon–oxygen double bond in structure VI. As the examples of CO2 and CH4 demonstrate, carbon has a valency of four because it attains the neon electronic configuration when it exercises that valency.

As noted earlier, an example of a triple bond is supplied by the nitrogen molecule, N2 (structure VII):

The Lewis structure for nitrogen molecules is given as structure XIII:

The existence of covalent bonds is not confined to neutral molecules. Such bonds also hold together polyatomic ions, that is ions that comprise modifications of several atoms. For example, when sodium hydroxide dissolves in water it forms ions of Na+ and (OH). Chemists indicate this by writing:

NaOH(s) = Na+(aq) + (OH)(aq) i

This shows that sodium hydroxide is an ionic compound, Na+(OH), but the hydroxide ion (OH) it contains is held together by a covalent bond. We can put an (OH) ion together by bonding an O ion to a hydrogen atom.

✦ A neutral oxygen atom contains six electrons (2s2 2p4). How many outer electrons will an O ion have?

✧ Seven. In addition to the original six electrons, an additional one is needed to give the ion a single negative charge.

✦ Draw a Lewis structure for (OH).

✧ A single bond now gives both oxygen and hydrogen a noble gas electronic configuration, see structure XIV.

(The additional electron is represented by the circle in structure XIV.)

The negative charge is now part of the whole oxygen–hydrogen complex, so it is placed outside brackets drawn around the Lewis structure.

At this point, you could be forgiven for thinking that the bonding and formulae of all chemical compounds can be explained by assuming that the atoms tend to attain noble gas electronic configurations. This is very far from being the case. For instance, the theory cannot account for the existence of transition element compounds such as FeCl2 and CuO, or for sulphur trioxide SO3. Nevertheless it can be usefully applied to many substances, and we now emphasize this by extending it to a new type of material.

4.5 Extended covalent bond

Figure 5 The structure of diamond. Each carbon atom is surrounded by four others located at the corners of a tetrahedron. The length of the C–C bond is 1.54 × 10−10 m.

Our study of covalent substances has so far been confined to compounds or elements that are composed of discrete molecules. However not all covalent substances are built in this way.

This point can be demonstrated by diamond, one of the forms of the element carbon. Although a different kind of covalent substance, its structure can be understood using the principles of the previous section. Diamonds are held together by C–C covalent bonds. Each carbon atom will form four such bonds; carbon has a valency of four, and in forming four single covalent bonds, attains the electronic configuration of neon.

In diamond, each atom forms four C–C electron–pair bonds, which point towards the corners of a regular tetrahedron (a pyramid with triangular base and sides). A fragment of the diamond structure is shown in Figure 5.

The C–C bonds are very strong, and you will notice that they spread out symmetrically through the structure in all directions. It is therefore very difficult to break into the structure with, say, a chisel.

Diamond is the hardest substance known and has a melting temperature (under pressure) of over 3500 °C. This extremely high melting point, together with the fact that it contains no discrete molecules, distinguishes diamond from the covalent compounds we have met so far. i

The absence of discrete molecules means that the C–C bonds must be broken if the structure is to be dismembered, and this accounts for the high melting temperature. Because the C–C bonds consist of shared electron pairs, diamond is covalent, but it is not molecular covalent in the sense we described earlier. We could argue to the contrary only if we were prepared to concede that entire diamond crystals could be regarded as giant molecules, which is contrary to experimental evidence.

The element silicon occurs in the same group of the periodic table as carbon, and it provides another example of an extended covalent substance – silicon dioxide (SiO2) which resembles carbon dioxide (CO2) only in its formula. Carbon dioxide is a gas at room temperature and comprises 0.03% of air by volume. It consists of discrete molecules and its Lewis structure shows that it is held together by carbon–oxygen double bonds (O|C|O). The physical properties of SiO2 are, however, very different from those of CO2.

Figure 6 Simplified structure of quartz (silicon dioxide). The length of the Si–O bond is 1.61 × 10−10 m.

The structure of quartz, a form of silicon dioxide, is shown in Figure 6. Here, both the silicon and oxygen atoms have noble gas electronic configurations. Each silicon ‘unit’ forms four single (electron-pair) bonds directed towards the corners of a tetrahedron and attains the electronic configuration of argon; each oxygen forms two single bonds and attains the configuration of neon.

The overall ratio of silicon to oxygen atoms in silica is 1 : 2. The structure is a three–dimensional network of Si–O bonds: there are no discrete molecules. Like diamond, it is an extended substance and it is not molecular covalent. Silicon dioxide (SiO2) should be a solid, and if the Si–O bonds are strong, as they are, the solid should be hard and it should melt at a high temperature. In fact, silicon dioxide SiO2 (or silica) occurs as the common mineral quartz, which is very hard, and it has the expected high melting temperature (1710 °C).

This brief comparison of silicon dioxide and carbon dioxide has highlighted the differences between extended and molecular covalent substances and it has demonstrated how profoundly the physical properties of a compound are affected by the type of chemical bonding between its constituents.

4.6 Electronegativity

When the atoms of non–metallic elements from the extreme right of the periodic table combine, they form covalent substances. However, combination of the atoms of a metallic element from the extreme left of the table with atoms of a non–metallic element from the extreme right, yields an ionic compound. Consider, for example, the third period, which runs from sodium to argon:

Na Mg Al Si P S Cl Ar

When two chlorine atoms combine, covalent Cl2 molecules of chlorine gas are formed. However, chlorine atoms combine with sodium atoms to form a solid, the ionic compound Na+Cl. Let us look at this contrast from an electronic point of view.

In structure XV, the Lewis structure of the Cl2 molecule is juxtaposed with a representation of the ions in sodium chloride.

The chloride ion has the electronic configuration of argon, with eight outer electrons, and the electron transferred from the sodium atom is marked by a black dot. Following the convention earlier, the outer electrons are grouped in pairs.

Notice that in both of these structures, the formation of a chemical bond involves the formation of a new electron pair in the outer electron configuration of chlorine. However, in Cl2, because the two atoms are identical, the electron pair must be equally shared between the two atoms whereas in NaCl it resides completely on the resulting chloride ion. According to this picture, ionic and covalent bonding are simply different aspects of the same process, because both involve the formation of electron pairs: the difference between them lies only in the extent to which those electron pairs are shared between atoms.

Atoms towards the right of the periodic table form covalent compounds when they interact; atoms towards the left behave somewhat differently. To explain this, we need to introduce the idea of electronegativity.

Table 4 Typical values for the
electronegativities of ten elements.
Element Electronegativity
sodium (Na) 0.9
lithium (Li) 1.0
silicon (Si) 1.8
hydrogen (H) 2.1
carbon (C) 2.5
bromine (Br) 2.8
nitrogen (N) 3.0
chlorine (Cl) 3.0
oxygen (O) 3.5
fluorine (F) 4.0

The electronegativity of an element is the power of an atom of the element to attract electrons to itself when entering into chemical combination (see Table 4). i

In the Cl2 molecule, the two identical atoms have an equal appetite for electrons: their electronegativities are equal, so the electron pair is shared equally between the two atoms. Now consider sodium chloride.

✦ Which atom is the more electronegative, sodium or chlorine?

✧ Chlorine: in sodium chloride, the electron pair has been completely taken over by chlorine which forms a chloride ion. If you imagine that the sodium and chlorine atoms have competed for electrons, the chlorine atoms have won.

This argument suggests that the electronegativity of chlorine, near the end of the third period, is greater than that of sodium at the beginning.

A further hint that this is correct comes from the ionization energies of some elements (see Figure 1). If an element has a low ionization energy, then an electron is removed from its atom with relative ease, and this suggests that the electronegativity of the element should be comparatively low. By the converse argument, high electronegativities are associated with high ionization energies; the ionization energy of chlorine, like its electronegativity, is greater than the corresponding value for sodium.

Figure 1 First ionization energies of elements up to radon.

More advanced theoretical arguments suggest that electronegativity increases steadily across the third period from sodium to chlorine. Within the periodic table in general, electronegativity usually increases from left to right across a period, and decreases from the top to bottom of a group. The most electronegative elements are in the top right–hand corner of the table – the top of groups and the end of periods. Ignoring helium, neon and argon, which form no chemical compounds, the most electronegative elements are first, fluorine, followed by oxygen, chlorine and nitrogen.

These trends in electronegativity can help us to predict whether one substance is more likely to be ionic than another. When atoms of high but similar electronegativity from the right–hand side of the periodic table bond together, covalent substances are formed. However, when atoms of high electronegativity combine with atoms of low electronegativity at the extreme left of the periodic table, ionic compounds tend to be formed.

It remains to consider what happens when atoms of low but similar electronegativity on the left of the periodic table combine with each other. This is the topic of the next subsection.

4.7 Metallic bonding

When the atoms of elements on the left of the periodic table come together, metals are formed. Because all the atoms have low electronegativities, they are prepared to surrender electrons to other atoms, either by electron transfer or electron–pair sharing. However, because the atoms have low electronegativity, none of them will readily accept these electrons, either by forming a negative ion, or by accepting a share in electron–pair bonds. In consequence, a pool of free electrons is created, which is like hot money – they are passed quickly from hand to hand, and can find no permanent home! Within this pool resides an array of positive ions formed from the atoms that supplied it (Figure 7). Ultimately, each atom shares its bonding electron or electrons with every other atom in the crystal. This is a simple but very useful model of metallic bonding.

Consider a specific example. When sodium atoms combine with one another, sodium metal is formed. Each sodium atom contributes one electron to the pool of free electrons, so each atom forms an Na+ ion with the noble gas configuration of neon.

Figure 7 A simple model of metallic bonding: in metallic elements and alloys, an array of positive ions is steeped in a pool of negatively charged free electrons. An alloy is a combination of elements, usually metallic elements in various proportions.

According to this model, therefore, sodium metal consists of an array of Na+ ions immersed in a pool of free electrons. The metal is held together by the ‘free’ electrons. It is energetically more favourable for the ions to be in close proximity with the sea of delocalized electrons rather than discrete atoms.

Because there is only one bonding electron per metal atom in sodium, the bonding is quite weak, and the melting point is low (98 °C). With magnesium, the next element in the period, the ion with a noble gas configuration is Mg2+. This double positive charge now interacts with the negative charge of the two free electrons per metal atom. The bonding is therefore stronger, and the melting point is correspondingly higher (650 °C).

You will of course be familiar with the distinctive properties of metals such as iron, aluminium, copper, silver and tin. Metals often have a lustrous appearance, and are good conductors of heat and electricity. The free electrons, which roam throughout the metal structure, account for the electrical conductivity. When a voltage is applied across two points on a piece of metal, the electron motion becomes less random, there is an overall movement of electrons between the two points, and an electric current flows.

Study comment We have discussed so far some of the simple theories of chemical bonding. There are more complex models which give a more complete description of chemical behaviour. One such model involves the concept of electrons occupying orbitals.

4.8 Summary of Section 4

Many substances are formed by atoms bonding together in order to attain the highly stable electronic configurations of the noble gases. Three main types of chemical bonding have been discussed:

These three kinds of bonding give rise to the distinctive properties of ionic, covalent and metallic substances. Whether a substance is ionic, covalent, or metallic is strongly influenced by the electronegativity of the atoms of which it is composed. In general, electronegativity increases from left to right across periods of the periodic table, and decreases down groups from the top to the bottom. This variation means that the non–metallic elements on the extreme right of the periodic table have high electronegativities; metallic elements at the extreme left have low electronegativities.

When non–metallic elements combine with themselves, or with each other to form binary compounds, their high but similar electronegativities result in the formation of shared electron–pair bonds. In other words, they form covalent substances. In many cases these substances can be represented by Lewis structures in which each non–metallic atom attains the electronic configuration of the noble gas at the end of its period.

When non–metallic elements from the extreme right of the periodic table form binary compounds with metallic elements from the extreme left, the large difference in electronegativity tends to generate ionic compounds. Such compounds can often be represented as assemblies of ions in which the metallic element has formed a positive ion with the electronic configuration of the preceding noble gas, and the non–metallic element had formed a negative ion with the electronic configuration of the following noble gas.

When elements from the left of the periodic table combine with themselves, or with each other, their low electronegativities generate assemblies of positive ions within a pool of free electrons. These free electrons account for the metallic properties of such elements, and also of the substances that they form by combination which are called alloys.

Question T10

Consider the compounds IBr, CaCl2 and CaMg2. One is ionic, one is covalent, and one is metallic. Identify which is which, and then match each compound to one of the three descriptions (a) to (c) below:

(a) White solid that melts at 772 °C. It is a poor conductor of electricity in the solid state, but a good one when melted or dissolved in water.

(b) Brown–black solid that melts at 41 °C to give a liquid with low electrical conductivity.

(c) Silvery looking solid that melts at 720 °C. Whether solid or molten, it is an excellent conductor of electricity.

Answer T10

Calcium chloride (CaCl2) is ionic, iodine bromide (IBr) is covalent and CaMg2 is metallic.

The properties listed are characteristic of (a) an ionic substance, (b) a covalent substance, and (c) a metallic substance. CaCl2 is a combination of elements from the extreme left and extreme right of the periodic table, so the electronegativity difference will be large and this will be the ionic compound. IBr will be covalent because it is a combination of elements of high electronegativity from the extreme right of the table. CaMg2 will be a metallic alloy because it is a combination of metallic elements with low electronegativity from the left of the table.

5 Closing items

5.1 Module summary

1

The first ionization energy of an atom is the energy required to remove its least tightly bound electron. When these values are plotted against atomic number, they trace a regular pattern. At the peaks are the noble (or inert) gases; just after the peaks are the very reactive alkali metals; and just before the peaks are the halogens, also very reactive.

2

The observed periodicity in the chemical properties of the elements enables them to be arranged in the form of a periodic table. The modern form of the table can be understood in terms of the electronic configuration of atoms, that is the distribution of the electrons across the various shells and subshells that they may occupy. No s subshell can contain more than two electrons, no p subshell more than six, no d subshell more than ten and no f subshell more than 14.

3

The periodic table is arranged so that all elements with similar outer electronic configurations (and hence similar chemical properties) appear in vertical columns, known as groups, in which the principal quantum number of the outer shell increases down the column. The horizontal rows in the periodic table, known as periods, typically contain several series of elements, each of which corresponds to the filling of a specific subshell.

4

At each end of the periodic table are groups of typical or main group elements. In between the main group elements are three series of transition elementstransition elements (corresponding to the filling of d subshells) and the lanthanides and actinides (corresponding to the filling of f subshells).

5

Simple theories of chemical bonding assume that the electronic configurations of the noble gases are especially stable, and that many elements tend to attain noble gas electronic configurations in their chemical reactions.

6

Valency is a property of an element that describes its ability to combine with other elements; it is helpful in predicting the formulae of certain compounds.

7

Subsection 4.3Ionic bonding involves the transfer of electrons from one atom to another, forming ions which have the outer electronic configuration of a noble gas. Ionic compounds consist of ions which in the solid state are held together by strong electrostatic forces; ionic compounds therefore tend to have comparatively high melting temperatures.

8

Subsection 4.4Covalent bonding involves the sharing of electrons between atoms so that the atoms each attain the outer electronic configuration of a noble gas. Covalent compounds are often made up of discrete molecules; the forces between these are in general less than the forces between ions, so the melting temperatures of covalent compounds tend to be less than those of ionic compounds. Subsection 4.5Extended covalent bonding, in which there are no discrete molecules, also arises.

9

The concept of Subsection 4.6electronegativity provides a link between ionic and covalent bonding. The electronegativity of an element is the power of an atom of the element to attract electrons to itself when entering into chemical combination.

10

Metallic bonding involves each atom sharing its bonding electron or electrons with every other atom in a crystal structure of positive ions.

5.2 Achievements

Having completed this module, you should be able to:

A1

Define the terms that are emboldened and flagged in the margins of the module.

A2

Outline the structure of the periodic table, and explain its relationship to the electronic configuration of atoms.

A3

Describe the main features and consequences of the electronic configuration of groups 0, 1 and 7 of the periodic table, together with those of the transition elements, the lanthanides and the actinides.

A4

Use the concept of valency to predict the structure of simple chemical compounds.

A5

Describe the characteristics of ionic, molecular covalent, extended covalent, and metallic bonding.

A6

Explain the relationship between the nature of the bonding in a compound and the properties of that compound.

A7

Draw Lewis structures for simple covalent compounds.

A8

Explain how the concept of electronegativity determines whether two atoms will combine to form an ionic, covalent or metallic compound.

Study comment You may now wish to take the Exit test for this module which tests these Achievements. If you prefer to study the module further before taking this test then return to the topModule contents to review some of the topics.

5.3 Exit test

Study comment Having completed this module, you should be able to answer the following questions each of which tests one or more of the Achievements.

Figure 4 Order of ascending energy for the subshells.

Question E1 (A2 and A3)

Radon, atomic number 86, marks the point in Figure 4 where the 6p level has become full. Use Figure 4 to predict the electronic configurations of the unknown elements 114, 118 and 119. (Use the symbol [Rn] for the electronic configuration of radon 86).

Which, if any, of these three elements is (a) a typical element, (b) a transition element, (c) an alkali metal, (d) a noble gas, (e) an actinide?

Answer E1

Element 114: [Rn] 7s2 5f14 6d10 7p2; element 118: [Rn] 7s2 5f14 6d10 7p6; element 119: [Rn] 7s2 5f14 6d10 7p6 8s1. Elements 110–120 have atoms with electronic configurations within the boundaries of Figure 4. After element 86, radon, where the 6p level is full, the order of filling is 7s, 5f, 6d, 7p and 8s.

The answers to the final part of the question are:

(a) typical elements: 114, 118, 119;

(b) transition element: none;

(c) alkali metal: 119;

(d) noble gas: 118;

(e) actinide: none.

(Reread Subsections 3.1 and 3.2 if you had difficulty with this question.)

Question E2 (A4)

(a) Aluminium has valency three. What are the chemical formulae of the simplest chemical compounds that it forms with the following elements:

(i) nitrogen, which has valency 3;

(ii) oxygen, which has valency 2;

(iii) carbon, which has valency 4.

(b) What are the valencies of the noble gases? Explain your answer.

Answer E2

(a)

(i) AlN;

(ii) Al2O3;

(iii) Al4C3.

These answers can be readily obtained by drawing the appropriate valency diagrams.

(b) Each noble gas has valency zero because each of them does not combine with other elements.

(Reread Subsections 4.1 and 4.2 if you had difficulty with this question.)

Question E3 (A2, A5 and A6)

An element Z is a solid that is a poor conductor of electricity. Z forms two fluorides with the formulae ZF4 and ZF6. The compound ZF6 is a colourless gas.

(a) In which group of the periodic table does Z lie? Explain how you deduce your answer from the properties described. What can you say about the electronic configuration of Z?

(b) What would you predict for the formula of the highest oxide of Z? Would you expect this oxide to be an ionic or covalent substance? Explain your reasoning.

(c) Deduce, as far as possible, the range of candidates for the identity of element Z.

Answer E3

(a) The fact that Z is a poor conductor of electricity tells us that it is a non-metal. The gaseous nature of ZF6 suggests that ZF6 is not ionic and that Z is therefore not a metallic element. The highest fluoride, ZF6 places Z in group 6 of the periodic table.

The electronic configuration of Z is therefore [A] s2 p4, where [A] is the electronic configuration of the noble gas of the highest atomic number that is lower than that of Z.

(b) The highest oxide of an element in group 6 would have the formula ZO3. The oxide ZO3 would be expected to be covalent because the difference in electronegativity between Z and O, both group 6 elements and both non-metals, is likely to be low.

(c) The non–metallic elements in group 6 are O, S and Se. The most likely candidates are S and Se. Oxygen can be ruled out because it is not a solid; its highest hydride is H2O and chloride is Cl2O.

(Reread Subsections 2.1, 2.2 and 4.6 if you had difficulty with this question.)

Question E4 (A4 and A7)

Nitrogen forms the compounds nitrosyl chloride (NOCl) and hydrazine (N2H4), which consist of discrete molecules. For each compound, draw diagrams of the molecules in which bonds are represented by lines, and each kind of atom has its common valency: nitrogen (3), oxygen (2), chlorine (1) and hydrogen (1). Then, for each case, draw the corresponding Lewis structure.

Answer E4

The only diagrams that give nitrogen a valency of three, oxygen a valency of two, and hydrogen and chlorine a valency of one are:

These diagrams also lead to Lewis structures in which each atom gains a noble gas electronic configuration with eight outer electrons.

(Reread Subsections 4.1 and 4.4 if you had difficulty with this question.)

Question E5 (A8)

Consider the following pairs of fluorides: (a) AlF3 and PF3; (b) CF4 and SnF4. One member of each pair is classified as an ionic substance, the other as a covalent substance. State which is which, expressing your reasons in terms of the electronegativities of the elements concerned.

Answer E5

AlF3 and SnF4 are ionic substances; CF4 and PF3 are covalent. Phosphorus lies to the right of aluminium in the third period of the periodic table, so it is more electronegative. The electronegativity of phosphorus is therefore closer to that of the very electronegative element fluorine. Thus the electronegativity difference between the elements is larger in AlF3 than in PF3, and AlF3 is more likely to be ionic. Likewise, tin lies below carbon in group 4, and so is less electronegative: this time, the electronegativity difference is larger in SnF4 than in CF4.

These conclusions are empirically correct: for example, CF4 and PF3 are gases at room temperature; AlF3 and SnF4 are solids that do not melt below 700 °C.

(Reread Subsection 4.6 if you had difficulty with this question.)

Study comment This is the final Exit test question. When you have completed the Exit test go back and and try the Subsection 1.2Fast track questions if you have not already done so.

If you have completed both the Fast track questions and the Exit test, then you have finished the module and may leave it here.

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