Concise Inorganic Chemistry (1991) PDF

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: · ·.·0. '.DRS. SOEOIJO ·.. : ' '. , :rJOKRODIHARJO DIP ::.'. ' PEit.UM: !KIP 15-KETlt.... SURABAYA. ·· · ·. CONCISE INORGANIC CHEMISTRY.. ' FOURTH EDITION __ J.D. Lee __ ~ Senior Lecturer in Inorganic Chemistr _ _M~'j I Loughborough UniversityofTechn~~~.J t ". CHAPMAN & HALL Universi1y and Professional Division London · New York · Tokyo · Melbourne · Madras UK Chapman & Hall, 2-6 Boundary Row, London SEI 8HN USA Chapman & Hall, 29 West 35th Street, New York NYIOOOI · JAPAN Chapman & Hall Japan. Thomson Publishing Japan, Hirakaw~cho Nemoto Building, 7F, 1-7-11 Hirnkawa-cho, Chiyoda-ku, Tokyo 102 AUSTRALIA Chapman & Hall Australia, Thomas Nelson Australia. 102 Dodds Street, South Melbourne, Victoria 3205 INDIA Chapman & Hi11l India. R. Seshadri. :12 Sernnd Main Road, CIT East, Mt I I 10 F \I I ~ ~ I Ne \ x I \ I I \ \ / I \ \ I I \ \ I I \ '-._X_.../..... -x..- I \ '-x-_... I 'x,,'-.x.-"',*' x 'x..'-x-,....;xI x / >'. Figure 1.2 Structures of the elements lithium to neon. 'x.. __ x" -x- -x- nitrogen, oxygen, fluorine and neon. Each has one more proton in the nucleus than. the preceding element, and the extra electrons go into a second orbit (Figure 1.2). This orbit is then full. In the next eight elements (with atomic numbers 11 to 18), the additiOnal electrons enter a third shell. The negatively charged electrons are attracted to the positive nucleus by electrostatic attraction. An electron near the nucleus is strongly attracted by the nucleus and has a low potential energy. An electron distant from the nucleus is less firmly held and has a high potential energy. ATOMIC SPECTRA OF HYDROGEN AND THE BOHR THEORY When atoms are heated or subjected to an electric discharge, they absorb energy, which is subsequently emitted ~s radiation. For example, if sodium chloride is heated in the flame of a aunsen burner' sodium atoms are produced which give rise to the characteristic yellow flame coloration. (There are two lines in the emission spectrum of sodium corresponding to wavelengths of589.0nm and 589.6nm.) Spectroscopy is a study of either the radiation absorbed or the radiation emitted. Atomic spectroscopy is an important technique for studying the energy and the arrangement of electrons in atoms. If a discharge is passed through hydrogen gas (H 2 ) at a low pressure, some hydrogen atoms (H) are formed, which emit light in the visible region. This light can he studied with u spectrometer, und is found to comprise a series of lines of different wavelengths. Four lines can be seen by eye, but many more are observed photographically in the ultraviolet region. The lines become increasingly close together as the wavelength (A.) decreases, until the continuum is reached (Figure 1.3). Wavelengths, in metres, are related to the frequency, v, in Hertz (cycles/second) by the equation: c v =- A  _ _ _ _A_T_O_M_I_C_S_P_EC_T_R_A_O_F_H_Y_D_R_O_G_E_N_A_N_D_T_H_E_B_O_H_R_T_H_E_O_R_Y~--- j [I) o< o< o< o< Cl) C') ;.. IO 0 r:- c'I Continuum co in co co CX) 'o:t 'o:t M 'o:t... 0 'o:t \ I H.. I I H13. Hy 1111111 Ha Hm Energy Figure 1.3 Spectrum of hydrogen in the visible region (Bahner series.) where c is the velocity of light (2.9979 x 108 ms- 1). In spectroscopy, frequencies are generally expressed as wave nurtlbets v, where v == 1/A.m ,. -. In 1885 Balmer showed that the wave number v of any line in the visible spectrum of atomic hydrogen could be given by the simple empirical formula: v = R(..!_ 22 -...!...) n2 where R is the. Rydberg constant and n has the values 3, 4. 5... thus giving a series of lines. The lines observed in the visible region are called the Balmer series. but several other series of lines may be observed in different regions of the spectrum (Table 1.1). Similar equations were found to hold for the lines in the other series in the hydrogen spectrum.. Lyman (l. - 1-) v = R 12 ll2 n = 2, 3, 4, 5... Ba Ith er v=R(l-1..) 22 n.2 n=3,4,5,6... Table 1.1 Spectral series found in atornic hydrogen Region of spectrum Lyman series ultraviolet Baitner series visible/ultraviolet Pascheil series infrared Brackett series infrared Pfund series infrared Humphries series infrared 0 [ _______A_T_O_M_I_C_S_T_R_U_C_T_U_R_E_A_N_D_T_H_E_PE_R_I_O_D_IC_T_A_B_L_E_ _ _ _ _ _~ Pasch en v= RG n\) n= 2 - 4, 5, 6, 7... Brackett v=R(~-_;) 4- rr 11=5,6,7,8... Pfund v=R(_!_ _. 52 _!_) n2 II = 6, 7, 8, 9... In the early years of this century, attempts were made to obtain a physical picture of the atom from this and other evidence. Thomson had shown in 1896 that the application of a high electrical potential across a gas gave electrons, suggesting that these were present in atoms. Rutherford suggested from alpha particle scattering experiments that a n atom con- sisted of a heavy positively charged nucleus with a sufficient number of electrons round it to make the atom electrically neutral. In 1913, Niels Bohr combined these ideas and suggested that the atomic nucleus was surrounded by electrons moving in orbits like planets round the sun. He was awarded the Nobel Prize for Physics in 1922 for his work on the structure of the atom. Several problems arise with this concept : I. The electrons might be expected to slow down gradually. 2. Why should electrons move in an orbit round t~e nucleus? · 3. Since the nucleus and electrons have opposite charges, they should attract each other. Thus one would expect the electrons to spiral inwards until eventually they collide with the nucleus. To explain these problems Bohr postulated: l. An electron did not radiate energy if it stayed in one orbit, and there- fore did not slow down. · 2. When an electron moved from one orbit to another it either radiated or absorbed energy. If it moved towards the nucleus energy was radiated and if it moved away from the nucleus energy was absorbed. 3. For an electron to remain in its orbit the electrostatic attraction between the electron and the nucleus which tends to pull the electron towards the nucleus must be equal to the centrifugat force which tends to throw the electron out of its orbit. For an electron of mass m, moving with a velocity v in an orbit of radius r ,.. mv- centnfugal force = -- r If the charge on the electron is e. the number of charges on the nucleus Z. and the permittivity of a vacuum £0.. f Cou 1om b1c attractive orce = -4-ze2 -, Jl£or so c _ __ _ _ _ A_T_O_M_IC_S_P_E_C_T_RA~O_F_H_Y~D_R_O_G_EN_A_N_D_T_H_E_B_O_H_R_T_H_E_O_R_Y_ _ _ _~j[Z] mv 2 Ze 2 -,- = 4neor 2 (1.1) hence Ze 2 v2 =--- (1.2) 4nE0mr According to Planck's quantum theory, energy is not continuous but is discrete. This means that energy occurs in 'packets' called quanta, of magnitude h/2rt, where h is Planck's constant. The energy of an electron in an orbit, that is its angular momentum mvr, must be equal to a whole number n of quanta. nh mvr=- 2n nh v=-. 2rtmr nih2 v2 = -,,.---~ 4rc2m2,2 Combining this with equation (1.2) Ze 2 n 2 h2 4ne0 mr - 4n 2 m 2 r2 hence t £ n2h2 = _(_l __ (1.3) rcme 2 Z For hydrogen the charge on the nucleus Z = I, and if n = I this gives a value r = 12 x 0.0529 nm n= 2 r = 22 x 0.0529 nm n=3 r = 32 x 0.0529 tlin This gives a picture of the hydrogen atom where an electron moves in circular orbits of radius proportional to 12 , 22 , 32 The atom will only radiate energy when the elettrOh jumps from one orbit to a.n other. The kinetic energy of an electron is -!mv 2 Rearranging equation (1.1) I 2 Ze2 E = -2mv = --· -- 81tEor Substituting for r using equation (1.3) Z 2e4 m E= 8e?,n 2 h2 If an electron jumps from an initial orbit i to a final orbit f. the change in energy D. E is [I][- ATOMIC STRUCTURE AND THE PERIODIC TABLE ------~=_] Zem( 1 I) 2 2 = 8£[,h ni - nr 2 Energy is related to wavelength {£ = hc:V so this equation is of the same form as the Rydberg equation: - ( 1.4) (Rydberg equation) Thus the Rydberg constant Z 2 e4 m R=. 8£.i1h 3 c The experimental value of R is 1.097373 x 107 m - 1, in good agreement with the theoretical value of 1.096776 x 107 m- 1 The Bohr theory provides an explanation of the atomic spectra of hydrogen. The different series of spectral lines can be obtained by varying the values of ni and fir in equation (1.4). Thus with nr = 1 and ni = 2, 3, 4... we obtain the Lyman series of lines in the UV region. With nr = 2 and ni = 3, 4, 5... we get the Balmer series of lines in the visible spectrum. Similarly, nr = 3 and n; = 4, 5, 6... gives the Paschen series, nr = 4 and ni = 5, 6, 7... gives the Brackett series, and nr = 6 and n;- = 7, 8, 9... gives the Pfund series. The various transitions which are possible between orbits are shown in Figure 1.4. REFINEMENTS TO THE BOHR THEORY It has been assumed that the nucleus remains stationary except for rotating on its own axis. This would be true if the mass of the nucleus were infinite, but the ratio of the mass of an electron to the mass of the hydrogen nucleus is 111836. The nucleus actually oscillates slightly about the centre of gravity, and to allow for this the mass of the electron m is replaced by the. reduced mass µ in equation (1.4): mM µ= m. + M where M is the mass of the nucleus. The inclusion of the mass of the nucleus explains why different isotopes of an element produce lines in the spectrum at slightly different wavenumbers. The orbits are sometimes denoted by the letters K, L, M, N... counting outwards from the nucleus, and they are also numbered 1, 2, 3, 4... This number is called the principal quantum number, which is given the symbol  ~~~~~~-R_E_F_IN_E_M~EN_T_S_T_O~T_H_E_B_O_H_·R~T_H_E_O_R_Y~~~~~~~~J[.2] 0 p Figure 1.4 Bohr orbits of hydrogen and the various series of spectral lines. n. It is therefore possible t define which circular orbit is under consideration by specifying the principal quantum number. When an electron moves from one orbit to another it should give a single sharp line in the spectrum, corresponding precisely to the energy dif- ference between the initial and final orbits. If the hydrogen spectrum is observe~ with a high resolution spectrometer it is found that some of the lines reveal 'fine structure'. This means that a line is really composed of several lines close together. Sommerfeld explained this splitting of lines by assuming that some of the orbits were eliiptical; and that they precessed in space round the nucleus. For the orbit dosest to the nucleus, the principal quantum number n = 1, and there is a circular orbit. For the next orbit, #1e principal quantum number n = 2, and both circular and elliptical orbits '~re possible. To define an elliptical orbit, a second quantum number k is needed. The shape of the ellipse is defined by the ratio Of the lengths of the major arid lninor axes. Thus major ·axis n minor axis = k k is called the azimuthal or subsidiary quantum 11utnber, and may have values from 1, 2... Ii. Thus for fl = 2, nlk may have the values 2/2 (circular  [}§] L _ ____A_T_O_M.I_C._:>_~~0~~~ -~N.D !_HE_P_E~!_?_P._l_S:__TABLE ----:--] orbit) and 2/1 (elliptical orbit). For the principal quantum number n = 3, nlk may have values 3/3 (circular), 3/2 (ellipse) and 3/1 (narrower ellipse). The presence of these extra orbits, which have slightly different energies k = 1 from each other , accounts for the extra lines in the spectrum revealed under high resolution. The original quantum number k has now been replaced by a new quantum number /, where I = k - 1. Thus for n=l l=O.S Bohr-Sommerfield 1en n = 3. ti = 2 I= 0 or 1 n=3 I= 0 or 1 or 2 II = 4 I = Oor 1 or 2 or 3 This explained why some of the spectral lines are split into two, three , four or more lines. In addition some spectral lines are split still further into two lines (a doubl-e t). This is explained by assuming that an electron spins on its axis in either a clockwise or an anticlockwise direction. Energy is quantized. and the vafue of the spin angular momentum was first con- sidered to be ms· h/2n, where ms is the spin quantum number with values of ±~- (Quantum mechanics has since shown the exact expression to be Vs(s + I)· h/2n, where s is either the spin quantum number or the resultant of several spins.) Zeeman showed that if atoms were placed in a strong magnetic field additional lines appeared on the spectrum. This is because elliptical orbits can only take up certain orientations with respect to the external field. rather than precessing freeiy. Each of these orientations is associated with a fourth quantum number m which can have values of l. (/ - I).... 0... (-/ + I).-/. Thus a single line in the normal spectrum will appear as (2/ + 1) lines if a magnetic field is applied. Thus in order to explain the spectrum· of the hydrogen atom, four quantum numbers are needed, as shown in Table 1.2. The spectra of other atoms may be explained in a similar manner. THE DUAL NATURE OF ELECTRONS-PARTICLES OR WAVES Symbol Values Principal quantum nurnl>er n l , 2. 3... Azimuthal or subsidiary quantum number 0, I,... (n - l) Magnetic quantum number m -1.... 0.... +l Spin quantum number ±!  THE HEISENBERG UNCERTAINTY PRINCIPLE.j ill] certain orbits. The electron is thus considered as a particle. In the 1920s it was shown that moving particles such as electrons behaved in some ways as waves. This is an important concept in explaining the electronic structure of atoms.. For some time light has been considered as either particles or waves. Certain materials such as potassium emit electrons when irradiated with visible light, or in sorne cases with ultraviolet light. This is called the photoelectric effect. It is explained by light travelling as particles called photons. If a photon collides with an electron, it can transfer its energy to the electron. If the energy of the photon is sufficiently large it can remove the electron from the surface of the metal. However, the phenomena of diffraction and interference of light can only be explained by assuming that light behaves as waves. In 1924, de Brogie postulated that the same dual character existed with electrons - sometimes they are considered as particles, and at other times it is more convenient to consider them as waves. Experimental evidence for the wave nature of electrons was obtained when diffraction rings were observed photographically when a stream of elec;trons was passed through a thin metal foil. Electron dif- fraction has now become a useful tool in determining molecular structure, particularly of gases. Wave mechanics is a means of studying the build-up of electron shells in atoms, and the shape of orbitals occupied by the electrons. THE HEISENBERG UNCERTAINTY PRINCIPLE Calculations on the Bohr model of an atom require precise information about the position of an electron an(f its velocity. It is difficult to measure both quantities accurately at the same time. An electron is too small to see and may only be observed if perturbed. For example, we could hit the electron with another particle such as a photon or an electton, or we could apply an electric or magnetic force to the electron. This will inevitably change the position of the electron, or its velocity and direction. Heisen- berg stated that the more precisely we cart define the position of an electron. the less certainly we are able to define its velocity. and vice versa. If l:!..x is the uncertainty in defining the position and !:!.. v the uncertainty in the velocity, the uncertainty principle may be expressed mathematically as:. h 6.x.6v~­. 4n where h = Planck's c_onstant = 6.6262 x 10- 34 J s. This implies that it is impossible to know both the position and the velocity exactly. The concept of an electron following a definite orbit, where its position and velocity are known exactly, must therefore be replaced by the prob- ability of finding an electron in a particular position, or in a particular volume of space. The Schrodinger wave equation provides a satisfactory description of an atom in these terms. Solutions to the wave equation are  ______A_T_O_M_I_C_S_T_R_U_C_T_U_R_E_A_N,_D_T_H_E__PE_R_I_O_D_IC TABLE 12 I L~ caJled wave functions and given the symbol tp. The probability of finding an electron at a point in space wh'ose coordinates are x, y and z is tp 2 (x, y, z). THE SCHRODINGER WA VE EQUATION For a standing wave (such as a vibrating string) of wavelength A., whose amplitude at any point along x may be described by a function f (x). it can be shown that: d2f(x) =- 4:n2 f(x) dx 2 A.2. If an electron is considered as a wave which moves in only one dimension then: An electron may move in three directions x. y and z so this becomes azw &2tp a2tp 4:n2.ax 2 + &y 2 + ()z 1 = -'")!' 'ljJ Using the · symbol V instead of the three partial differentials, this is shortened to The de Broglie relationship states tha.t A.=..!!.. mv (where h is Planck's constant, m is the mass of an electron and v its velocity); hence: or (1.5) However, the total energy of the system Eis made up of the kinetic energy K plus the potential energy V E = K+ V. so K=E-V But the kinetic energy = ~mv 2 so  [ __________ TH_E_SC_H_R_o_ · D_IN_G_E_R_W_A_V_E_EQ_UA_T_IO_N_ _ _ _ _ _ _ _ J [ill !mv 2 = E - V and v 2 = -2 (E - V) m Substituting for v 2 in equation (1.5) gives the well-known form of the Schrodinger equation 8n 2 m 2 V tjJ +IT° (E - V)tjJ =0 Acceptable solutions to the wave equation, that is solutions which are physically possible, rriust have certain properties: 1. tjJ must be continuous. 2. tjJ must be fi.nite. 3. tjJ must be single valued. 4. The probability of finding the electron over all the space from plusinfinity to minus infinity must be equal to one. The probability of finding an electron at a point x, y, z is tjJ 2, so +oo J"1 -oo 2 dxdydz =1 Several wave functions called 'i'i. '\jJ 2 , '\jJ3.. will satisfy these conditions to the wave equation, and each of these has a corresponding energy E1,.£2 , E3.. Each of these wave functions 'ljJ" '\jJ 2 , etc. is called an orbital, by analogy with the orbits in the Bohr theory. In a hydrogen atom, the single electron normally occupies the lowest bf the energy levels E 1 This is called the ground state. The corresponding wave function 'lj/ 1 describes the orbital, that is the volume in space where there is a high probability of finding the electron. For a given type of atom, there are a number of solutions to the wave equation which are acceptable, and each orbital may be described uniquely by a set of three quantum numbers, n, I and m. (These are t_h e same quantum numbers - principal, subsidiary and magnetic - as were used in the Bohr theory). The subsidiary quantum number I describes the shape of the orbital occupied by the electron. l may have values 0, 1, 2 or 3. When I= 0, the orbital is spherical and is called an s orbital; when I = l, the orbital is dumb-bell shaped and is called a p orbital; when I= 2, the orbital is double dumb-bell shaped and is called a d orbital; and when I = 3 a more complicated f orbital is formed (see Figure 1.6). the letters s, p, d and i come from the spectroscopic terms sharp, principal, diffuse and.funda- mental, which were used to describe the lines in the atomic spectra. Examination of a list of all the allowed solutions to the wave equation shows that the orbitals fall into groups. In the first group of solutions the value of the wave function 'ljJ, and  ---------------- ATOMIC STRUCTURE AND THE PERIODIC TABLE y y y 2s y Figure 1.6 (a) Wave functions 'ljl for s. p and ,1 atomic orbitals: (i) s orbital. 2s; (ii) p 1---~)( orbital, 2px; (iii) d orbital, 3d.n.. (Note that the + and - signs refer to symmetry, not charge.) (b) Different ways of repre~enting 'ljl 2 for a 2p orbital (as a contour / diagram or as a 90% boundary surface). hence the probabiHty of finding the electron 1jl2 depends only on the distance r from the nycleus, and is the same in all directions. y \jl == f (r) This leads to a spherical orbital, and occurs when the subsidiary quantum number I is zero. These are called s orbitals. When I = 0, the magnetic quantum number m = 0, so there is only one such orbital for each value of n. In the second group of solutions to the wave equation, 'ljJ depends both on the distance from the nucleus, and on the direction in space (x, y or z). Orbitals of this kind occur when the subsidiary quantum number I = 1. These are called p orbitals and there are three possible values cif the magnetic quantum number (m = -1, 0and+1). There are therefore three orbitals which are identical in energy, sh.a pe and size, which differ only in their direction in space.These three solutions to the wave equation may be written (a) ip-1" = f(r). f(x) 'll'.v = f(r). f(y) "Pz = f(r). f(z) Orbitals that are identical in energy are termed degenerate, and thus three degenerate p orbitals occur for each of tQe values of n = 2, 3, 4... The third group of solutions to the wave equation depend on the  RADIAL AND ANGULAR FUNCTIONS I [ill Table 1.3 Atomic orbitals Principal Subsidiary Magnetic Symbol quantum quantum quantum number number numbers n I m 0 0 ls (one orbital) 2 0 0 2s (one orbital) 2 1 -1, 0, +1 2p (three orbitals) 3 0 0 3s (one orbital) 3 1 -1, 0, +1 3p (three orbitals) 3 2 -2, -1, 0, +1, +2 3d (five orbitals) 4 0 0 4s (one orbital) 4 1 -1,0, +1 4p (three orbitals) 4 2 -2, -1, 0, +1, +2 4d (five orbitals) 4 3 -3, -2, -1, 0, +l, +2, -3 4/ (seven orbitals) distance from the nucleus r and also on two directions in space, for example 'fl = f(r). f(x). f(y) This group of orbitals has I = 2, atld these are called d orbitals. there are five solutions corresponding tom = -2, -1, 0, + 1 and +2, and these are all equal in energy. Thus five degenerate d orbitals occur for each of the values of n = 3, 4, 5.... A further set of solutions occurs whetl I = 3, and- these are called f orbitals. There are seven values of tn: -3, -2, -l, O", +l, +2 and +3, and seven degenerate f orbitals are formed when n = 4, 5, 6.... RADIAL AND ANGULAR FUNCTIONS The Schrodinger equation can be solved completely for the hydrogen atom, and for related ions which have only one electron such as He+ and u 2 +. For other atoms only approx:irnate solutions can be obtained. For most calculations, it is simpler to solve the wave equation if the cartesian coordinates x, y and z are converted into polar coordinates r, e and. The coordinates of the point A measured from the origin are x, y, a.rid z in e cartesian coordinates, and r, and in polar coordinates. It c·an be seen that the two sets of coorditlates are related by the following expressions: z =,cos a y= r sin e sin x' = r sin e cos ¢' The SchrOdinger equation is usually written: 8n 2m. · ·· V 2ip + h2 (£.- V)\j! = 0 QI] j~~~~~~~A_T_O_M~IC_S_T_R_U_C_T_U_RE~AN~D_T_H_E_P_E_R_IO_D_I_C_T_A_B_LE~~~~~~-' z ,,"" ,, ------ ------ / r----- z--- 1 I I I I I I I I I I y x Figure 1.7 The relationship between cartesian and polar coordinates. where 2 o2'ljl iP'ljJ 8 2'\jl \7'1jJ=-+-+- 8x2 {)y2 ()z2 Changing to polar coordinates, \72'\jl becomes l.~(r2 o'ljl) + 1. a21" +~.~(sine o'ljl) r2 or or r 2 sin 2 e acp 2 2 r sine ae ae The solution of this is of the form ljl = R(r). 0(0), (cp) (1.6) R(r) is a function that depends on the distance from the nucleus, which in turn depends on the quantum numbers n and I 0(0) is a function of 0; which depends on the quantum numbers I and m (cj>) is a function of cp, which depends only on the quantum number m Equation (1.6) may be rewritten '11 = R(r)nr. Amr This splits the wave function into two.pa~ts which can be solved separately: 1. R(r) the radial function, which depends on the· quantum numbers n and/. 2. Amr the total angular wave function, which depends on the quantum numbers m and I. The radial function R has no physical meaning, but R 2 gives the probability of finding the electron in a small volume dv near the point at which R is measured. For a given value of r the number of small volumes is 4nr 2 , so the probability of the electron being at a distance. r from ·the nucleus is 4nr 2 R 2 This is called the radial distribution function. Graphs of the  RADIAL AND ANGULAR FUNCTIONS ·--~IT~ 1s 5 10 '.~ 0 5 10 2s ~'~~ a: 2·0l ~.:. 5 ,2 10. i·5. 1·0 ~ - 0·5~~~ 0 5 10 11·0·5r ~ 0·5 0 5 10. 1·5r 1·0. ~ · 0·5~3d 0 5 10 , 1.A1 Figure 1.8 Radial distribution functions for various orbitals in the hydrogen atom. radial distribution function for hydrogen plotted against r are shown in Figure 1.8. · These diagrams show that the probability is zero at the nucleus (as r = 0), and by examining the plots for ls, 2s and 3s that the most probable distance increases markedly as the principal quantum number increases. Furthermore, by comparing the plots for 2s and 2p, or 3s, 3p and 3d it can be seen that the most probable radius decreases slightly as the subsidi_ary quantum number increases. All the~· orbitals except the first one (ls) have a shell-like structure, rather like an onion or a hailstone, consisting of concentric layers of electron density. Similarly, all bun he first p orbitals (2p) and the first d orbitals (3d) have a shell structure. The angular function A depends only on the direction, and is indepen- dent of the distance from the nucleus (r). Thus A 2 is the probability of  z x 1s z z z x x y 2p. y 3dx2-y2 3dz2 Figure 1.9 801,mdary surface for the angular part of the wave function A (0, 4>) for the 2s, 2p and 3d orbitals for a hydrogen atom shown as polar diagrams.  L--~~~~~~~-R_A_D_I_A_L_A_N_D_A_N_G_U_L_A_R_F_U_N_C_T_IO_N~S~~~~~~~~---'' O_I] finding an electron at a given direction e, at any distance from the nucleus to infinity. The angular functions Aare plotted as polar diagrams in Figure 1.9. It must be emphasized that these polar diagrams do not represent the total wave function u, but only the angular part of the wave function. (The total wave function is made up from contributions from both the radial and the angular functions.) 'ljl = R(r). A Thus the probability of finding an electron simultaneously at a distance r and in a given direction e' is '\/';. 0 ,. 2 2 'ljl;,0. = R (r). A (e, ) Polar diagrams, that is drawings of the the angular part of the wave function, are commonly used to illustrate the overlap of orbitals giving bonding between atoms. Polar diagrams are quite good for this purpose, as they show the signs + and - relating to the symmetry of the angular function. For bonding like signs must overlap. These shapes are slightly different from the shapes of the total wave function. There are several points about such diagrams: l. It is difficult to picture an angular wave function as a mathematical equation. It is much easier to visualize a boundary surface , that is a solid shape. which for example contains 90% of the electron density. To emphasize that 'ljl is a continuous.function, the boundary surfaces have been extended up to the nucleus in Figure 1.9. For p orbitals the electron density is zero at the nucleus, and sotne texts show a p orbital a' two spheres which do not touch. · 2. These drawings shaw the symmetry for the ls, 2p , 3d orbitals. However, in the others, 2s, 3s, 4s... , 3p, 4p, Sp... , 4d, Sd... the sign (symmetry) changes inside the boundary surface Of the orbital. this is readily seen as nodes in the graphs Of the radial functions (Figure 1.8). z z z 2pz Figure 1.10 The angular part of the wave function squared A 2 (0, ) for the 2p orbitals for a hydrogen atom.  z x y 15 z x z x z x y y 2py y 2p,. 2Pz z z z x x x 3dyz z z x x 3d,2 Figure 1.11 Total wave function (orbitals) for hydrogen.  L--~~~~~~-B_U_IL_D_-U_P_O_F~T_H_E_E_LE_M_E_N_T_S_,_I-ru~N_D_'S_R~U_L_E~~~~~~'----'j G'IJ 3. The probability of finding an electron at a direction e, is· the wave function squared, A 2 , or more precisely 'ljJ~'ljJ~. The diagrams in Figure 1. 9 are of the angular part of the wave function A, not A 2 Squaring does not change the shape of an s orbital, but it elongates the lobes of p orbitals (Figure 1.10). Some books use elongated p orbitals, but strictly these should not have signs, as squaring removes · any sign from the symmetry. Despite this, many authors draw shapes approximating to the probabilities, i.e. squared wave functions, and put the signs of the wave function on the lobes, and refer to both the shapes and the wave functions as orbitals. 4. A full representation of the probability of finding an electron requires the total wave function squared and includes both the radial and angular probabilities squared. It really needs a three-dimensional model to display this probability, and show the shapes of the orbitals. It is difficult to do this adequately on a two-dimensional piece of paper, but a representation is shown in Figure 1.11. The orbitals are not drawn to scale. Note that the p orbitals are not simply two spheres, but are ellipsoids of revolution. Thus the 2px orbital is spherically symmetrical about the x axis, but is not spherical in the other direction. Similarly the Pr orbital is spherically symmetrical about the y axis, and both the Pz and the 3dz! are spherically symmetrical about the z axis.. PAULl EXCLUSION PRINCIPLE Three quantum :numbers n, I and m are needed to define an orbital. Each orbital may hold up to two electrons, provided they have opposite spihs. An extra ·quantum number is required to define the spin of an electron in an orbital. Thus four quantum numbers are needed to define the energy of an electron in an atom. The Pauli exclusion principle states that no two electrons in one atom can have all four quantum numbers the ~ame. Ry permutating the quantum numbers, the maximum number of electrons which can be contained in each main energy level can be calculated (see Figure 1.12). BUILD-UP OF THE ELEMENTS, HUND'S RULE When atoms are in their ground state, the electrons occupy the lowest possible energy levels. The simplest element, hydrogen, has one electron, which occupies the ls level; this level has the principal quantum number n = 1, and the subsidiary quantum number I = 0. Helium has two electrons. The second electron also occupies the lslevel. This is possible because the two electrons have opposite spins. This level is now full. The next atom lithium has three electrons. The third electron occupies the next lowest level. This is the 2s level, which has the principal quantum number n = 2 and subsidiary quantum number I = 0. Principal Quantum 1 2 3 No.n I I I Subsidiary Quantum No. / I 0 I 0 I 1 I 0 I 1 J Magnetic Quantum.No.m _J_ I 0 _L I ~ I _L -l I _J_ I 0 _L I +1 I I 0 _L I -1 I \ 0 I I +l I _L -2 I -1 I I l I +1 I +2 l :Spin Quantwn · No.m, I n n ,.. I I - n n n n 11 i i 11 I I I I I I +! -t +t -! +! -t +t -t +t -t. +! -t +! -t.+t -t +t -t +t -t +t -t +! -! +t -! +t -! '--y-J ~ '-v-' Total 2J 2J 6p 2.s 6p tOd electrons electrons electrons Total 8 electrons Total 18.electrons N 4 I I I I I 0 1 2 3 I 1 I., 1 1 ! 1 1 ,--·, 1 \ r T ·1 ,,, 0 -1 r1-, r1-, ,- r1-, , 0 1 +1 -2 -1 t-, r11 ri n n n n 0 +l +2 -3 m,+t -! +! -! +! -t +! -! +! -! +! -! +t -! +t -! +t -t +t -! +t -t +t -t +! -t +! -! +t -t +! -t '-v-' -2..______________ ~ -1 r-, 0. ---------------- 1. 1 r1i. ri r1-, 3 21 ~ ~ w Total 32 electrons Figure 1.12 Quantum numbers, the permissible number of electrons and the shape of the periodic table.  L-~~~~~~-B_U_IL_D_-U_P~O_F_T_H_E_E_L_EM~EN_T_S~,_H_U_N_D~'S_R_U_L_E~~~~~~----''!}Ij The fourth electron in beryllium also occupies the 2s level. Boron must have its fifth electron in the 2p level as the 2s level is full. The sixth electron in carbon is also in the 2p level. Hund's rule states. that the number of unpaired electrons in a given energy level is a maximum. Thus in the ground state the two p electrons in carbon are unpaired. They occupy separate p orbitals and have parallel spins. Similarly in nitrogen the three p electrons are unpaired and have parallel spins. To show the positions of the electrons in an atom, the symbols Is, 2s, 2p, etc. are used to denote the main energy level and sub-level. A superscript indicates the number of electrons in each set of orbitals. Thus for hydrogen, the Is orbital contains one electron, and this is shown as Is 1 For helium the ls orbital contains two electrotls, denoted ls 2. The electronic structures of the first few 'atoms in the periodic table tnay be written: H ls 1 He ls2 Li ls2 2sl Be ls2 2s2 B ls2 2s2 2pl c ls2 2s2 2p2 N ls2 2s2 2p3 0 ls2 2s2 2p4 F ls2 2s2 2p5 Ne ls2 2s2 2p6 Na· ls2 2s2 2p6 3s 1 An alternative way of shOwing the electroniC structure of an atom is.to draw boxes for orbitals, and arrows for the electrons. ls 2s 2p Electronic structure of.H atom in the ground state EJ D I I I I ls 2s 2p Electronic structure of He.atom in the ground state [!±] D IIlI ls 2s 2p Electronic structure of Li atom in the ground state [!±] EJ I I I I ls 2s 2p Electronic structure of Be atom in the ground state [!±] lliJ I ls 2s 2p Electronic structure of B atoln in the ground state ~ ~ It I I I  [EJ I ATOMIC STRUCTURE AND THE PERIODIC TABLE ls 2s 2p Electronic structure of C atom in the ground state [!±] [!±] Ir Ir I I ls 2s 2p Electronic structure of N atom in the ground state [!±] (!±] Ir It It I ls 25 2p Electronic structure of 0 atom in the ground state [!±] [!±] !Hit Ir I ls 2s 2p ~ (!±] lt.i.IHlr I Electronic structure of F atom in the ground state ls 25 2p Js Jp Electronic structure of Ne atom in the ground state 5J 5J It.i. IHIHI D I I I I ls 2.s 2p 3s Jp G±1 G±l lr.i.IHlr.i.I EJ Electronic structure of Na atom in the ground state I I I I The process continues in a similar way. SEQUENCE OF ENERGY LEVELS It is important to know the sequence in which the energy levels are filled. Figure 1.13 is a useful aid. From this it can be seen that the order of filling of energy levels is: ls, 2s, 2p, 3s, 3p, 4s, 3d, 4p, Ss, 4d, Sp, 6s, 4/, Sd, 6p, 7s, etc. After the ls, 2s, '{.p, 3s and 3p levels have been filled at argon, the next two electrons go into the 4s level. This gives the elements potassium and calcium. Once the 4s level is full the 3d level is the next lowest in energy, not the 3p level. Thus tl:ie 3d starts to fill at scandium. The elements from scandium to copper have two electrons in the 4s level and an incomplete 3d level, and all behave in a similar manner chemically. Such a series of atoms is known as a tninsition series.. · A second transition series starts after the 5s orbital has been tilled, at strontium , because in the next element, yttrium, the 4d level begins to fill up. A third transition series starts at- lanthanum where the electrons start to fill the Sd level after the 6d level has been filled with two electrons. A further complication arises here because after lanthanum, which has one electron in the 5d level, the 4/ level begins to fill, giving the elements from cerium to lutetium with from one to 14/ electrons. These are 3 Sequence of filling sometimes called the inner transition elements, but are usually known as ·els. the lanthanides or rare earth metals.  '----A_R_R_A_N_G_E_M_EN_T_O_F_T_H_E_E_L_E_M_E_N_T_S_I_N_G_R_O_U_P_S_I_N_T_H_E_P_E_R_IO_D_I_C_T_A_B_L_E_ ___,J J 25 j ARRANGEMENT OF THE ELEMENTS IN GROUPS IN THE PERIODIC TABLE The chemical properties of an element are largely governed by the number of electrons in the outer shell, and their arrangement. If the elements are arranged in groups which have the same outer electronic arrangement, then elements within a group should show · similarities in chemical · and physical properties. One great advantage of this is that initially it is only necessary to learn the properties of each group rather than the properties of each individual element. Elements with ones electron in their outer shell are called Group I (the alkali metals) and elements with two s electrons in their outer shell are called Group II (the alkaline earth metals). These two groups are known as the s-block elements, because their properties result from the presence of s electrons. Elements with three electrons in their outer shell (two s electrons and one p electron) are called Group III, and similarly Group IV elements have four outer electrons, Group V elements have five outer electrons, Group VI elements have six outer electrons and Group VII elements have seven outer electrons; Group 0 elements have a full outer shell of electrons so that the next shell is empty; hence the group name. Groups III, IV, V, VI, VII and 0 all have p orbitals filled and because their properties are dependent on the · presence of p electrons, they are called jointly the p-block elements... In a similarway, elements where d orbitals are being filled are called the d-block, or transition elements. In these, d electrons are being added to the penultimate shell. · · Finally, elements where f orbitals are filling are called the /-block, and here the f electrons are entering the antepenultimate (or second from the outside) shell. In the periodic tab!~ (Table 1.4), the elements are arranged in order of increasing atomic number, that is in order of increased nuclear charge, or increased number of orbital electrons. Thus each element contains one more orbital electron than the preceding element. Instead of listing the 103 elements as one long list, the periodic table arranges them into several horizontal rows or periods, in such a way that each row begins with an alkali metal and ends with a noble gas. The sequence in which the various energy levels are filled detertnines the number of elements in each period, and the periOdic table can be divided into fout main regiotls according to whether the s, p, d or f levels are bein~ filled. 1st period ls elements in this period 2 2nd period it 2p elements in this period 8 3rd period 3s 3p elements in this period 8 4th period 4s 3d 4p elemet1ts in this period 18 5th period 5s 4d Sp elements in this period 18 6th period 6s 4/ 5d 6p elements in this period 32 Table 1.4 The periodic table s-block p-block '~ p. I II Ill IV v VI VII 0 I IH IH 2He · 2 3Li 4Be SB "C 7N KO >F 111 Ne 3 uNa 12Mg d-block 13 AI 1,.Si ISp. "s 17 CI IKAr 4 '''K. 20 Ca 21 sc 22Ti nv 24Cr 2SMn 2"Fe i1co · 2KNi 211 Cu :11 zn 31 Ga J2Ge 33 As 34 Se 35 Br Jr.Kr 5 37 Rb ·-3Ksr J~y 4ozr 4'Nb 42Mo 4J-rc '4Ru 4SRh 41 'Pd 4'Ag 4KCd 4 '>1n sosn 51 Sb s2Te 531 S4Xe· 0 sscs S(;Ba 57La 72 Hf '3Ta 14w 7SRe '''Os 77 Ir 7Kpt 7'>Au ~'Hg KITI K2pb KJBi IWpO KS At ~'Rn 7 K7Fr 1111 Ra 1 "'Ac /-block. Lanttianides sxce S >pr '~ 1 Nd 1 ' 1Pm ''2Sm '' 3 Eu ""Gd ''5Tb '"'Dy "7Ho · '.xEr ''''Tm 70 Yb 71 Lu Actinides ioTh '"Pa >2u ''·'Np >4pu " 5Am Cm 97 Bk >KCf '"'Es 11 1 MFm 1111 Md 1112 No 1113 Lr  PROBLEMS The alkali metals appear in a vertical column labelled Group I, in which all elements have ones electron in their outer shell, an'd hence have similar properties. Thus w~en one element in a group reacts with a reagent, the other elements in the group will probably react similarly, forming com- pounds which have similar formulae. Thus reactibns of new compounds and their formulae may be predicted by analogy with kMwn compounds. Similarly the noble gases all appear in a vertical column labelled Group 0, and all have a complete outer shell of electrons. This is called the long form of the periodic table. It has many advantages, the most important being that it emphasizes the similarity of properties within a group and the relation between the group and the electron structure. The d-block elements are referred to as the transition elements as they are situated between the s- and p-blocks. Hydrogen and helium differ from the rest of the elements because there are no p orbitals in the first shell. Helium obviously belongs to Group 0, the noble gases, which are chemically inactive because their outer shell of electrons is full. Hydrogen is more difficult to place in a group. It could be included in Group I because it has one s electron in its outer shell, is univalent and commonly forms univalent positive ions.. However, hydrogen is not a metal and is a gas whilst Li, Na, K, Rb and Cs are metals and are solids. Similarly, hydrogen could be included in Group VII because it is one electron short of a complete shell, or in Group IV because its outer shell is half full. Hydrogen does not resemble the alkali metals, the halogens or Group IV very closely. Hydrogen atoms are extremely small, and have many unique properties. Thus there is a case for placing hydrogen in a group on its own. FURTHER READING Karplus, M. and Porter, R.N. (1971) Atoms and Molecules, Benjamin, New York. Greenwood, N.N. (1980) Principles of Atomic Orbitals, Royal Institute of Chemistry Monographs for teachers No. 8, 3rd ed., London. PROBLEMS 1. Name the first five series of lines that occur in the atomic spectrum of hydrogen. Indicate the region in the electromagnetic spectrum where these series occur, and give a general equation for the wavenumber applicable to all the series. 2. What are the assumptions on which the Bohr theory of the structure of the hydrogen atom is based? 3. Give the equation which explains the different series of lines in the atomic spectrum of hydrogen. Who is the equation named after? Explain the various terms involved. 4. (a) Calculate the radii of the first three Bohr orbits for hydrogen. (Planck's 6onstal'lt h = 6.6262 x 10- 34 Js; mass of electron 0J ~'______A_T_O_M_I_C_S_TR_U_CTURE AND THE PERIODIC TABL_E_ _ ,.___J m = 9.1091x10- 3 kg; charge on electron e = 1.60210 x 10- 1''C; permittiv.ity of vacuum Eo = 8.854185 x 10- 12 kg- 1 m- 3 A 2.) (Answers: 0.529 x 10- 10 m; 2.12 x 10- 10 m; 4.76 x 10- 10 m; that is 0.529A 2.12A and 4.76A.) · (b) Use these radii to calculate the velocity of an electron in each of these three orbits. · (Answers: 2.19 x 106 ms"" 1; 1.09 x 106 ms- ; 7.29 x 105 ms- 1.) 5. The Balmer series of spectral lines for hydrogen appear in the visible region. What is the lower energy level that these electronic transitions start from, and what transitions correspond to the spectral lines at 379.0 nm and 430.Q nm respectively? 6. What is the wavenumber and wavelength of the first transition in the Lyman, Balmer and Paschen series in the atomic spectra of hydrogen? 7. Which of the following species does the Bohr theory apply to? (a) H, (b) H+, (c) He, (d) He+, (e) Li, (f) u+, (g) u+ 2 , (h) Be, (g) Be+, (h) Be2 +, (i) BeJ+. 8. How does the Bohr theory of the hydrogen atom differ from that of Schrodinger? 9. (a) Write down the general form of the Schrodinger equation and define each of the terms in it, (b) Solutions to the wave equation that are physically possible must have four special properties. What are they? 10. What is a radial distribution function? Draw this function for.the ls; 2s, 3s, 2p, 3p and 4p orbitals in a hydrogen atom. 11. Explain (a) the Pauli exclusion principle, and (b) Hund's rule. Show how these are used to specify the electronic arrangements of the first 20 elements in the periodic table. 12. What is an orbital? Draw the shapes of the ls; 2s, 2px, 2py, 2pz, 3dxy 3dxz 3dyz 3dx2-y2 and 3dz2 orbitals. 13. Give the names and symbols of the four quantum numbers required to define the energy of electrons in atoms. What do these quantum numbers relate to, and what numerical values are possible for each? Show how the shape of the periodic table is related to these quantum numbers. 14. The first shell may contain up to 2 electrons, the second shell up to 8, the third shell up to 18, and the fourth shell up to 32. Explain this arrangement in terms of quantum numbers. 15. Give the values of the four quantum numbers for each electron in the ground state for (a) the oxygen atom, and (b) the scandium atom. (Use positive values for m 1 and ms first.) 16. Give the sequence in which the energy levels in an atom are filled with electrons. Write the electronic configurations for the elements of atomic number 6, 11, 17 and 25, and from this decide to which group in the periodic table each element belongs. 17. Give the name and symbol for each of the atoms which have the ground state electronic configurations in their outer shells: (a) 2s2 , (b) 3s2 3p 5 , (c) 3s 2 3p 6 4s 2 , (d) 3s 2 3p6 3d 6 4s 2 ; (e) 5s 2 5p 2 , (f) 5s 2 5p 6  Introduction to 2 bonding ATTAINMENT OF A STABLE CONFIGURATION How do atoms combine to form molecules and why do atoms form bonds? A molecule will only be formed if it is more stable, and has a lower energy, than the individual atoms.. To understand what is happening in terms of electronic structure, con- sider first the Group 0 elements. These comprise the noble gases, helium, neon. argon, krypton, xenon and radon, which are noteworthy for their chemical inertness. Atoms of the ·noble gases do not normally react with any other atoms, and their molecules are monatomic, i.e. contain only one atom. The lack of reactivity is because the atoms already have a low energy, and it cannot be lowered further by forming compounds. The low energy of the noble-gases is associated with their having a complete outer shell of electrons. This is often called a noble gas structure, and it is an exceptionally stable arrangement of electrons. Normally only electrons in the outermo~t shell of an atom are involved in forming bonds, and by forming bonds each atom acquires a stable electron configuration. The most stable electronic arrangement is a noble gas structure. and many molecules have lhis arrangement. However, less stable arrangements than this are commonly attained by transition elements·. TYPES OF BONDS Atoms may attain a stable electronic configuration in three different ways: by losing electrons, by gaining electrons, or by sharing electrons. Elements may be divided into: I. Electmpositive elements, whose atoms give up one or more electrons fairly readily. 2. electronegative elements. which will accept electrons. 3. Elements which have little tendency to Jose or gain electrons. Three different types of bond may be formed, depending on the electropositive or electronegative character of the atoms involved..._______T_R_A_N_S_IT_I_O_N_S_·B_E_lW_E_E_N_T_H_E_M_A_IN_T_Y_PE_S_o_ · _F_B_O__ N_D_IN_G _ _ _ _ ____J] QIJ ·Electropositive element + JIonic bond Electronegative element Electronegative element J + Covalent bond Electronegative element Electropositlve element } + Metallic bond Electropositive element Ionic bonding involves the complete transfer of one or more electrons from one atom to another. Covalent bonding involves the sharing of a pair of electrons between two atoms, and in metallic bonding the valency electrons are free to move throughout the whole crystal. These types of bonds 'a re idealized c>r extreme representations, and though one type generally predominates, in mosrsubstances the bond type is somewhere between these extreme forms. For example, lithium chloride is considered to be an ionic compound, but it is soluble in alcohol, which suggests that it also possesses a small amount of covalent character. If the three extreme bond types are placed at the corners of a triangle, then compounds with bonds predominantly of one type will be represented as points near the corners. Compounds with bonds intermediate between two types will occur along an edge of the triangle, whilst compounds with bonds showing some characteristics of all three types are shown as points inside the triangle. Metallic Li I\ Ag Na3 Bi I \ Sn Na3 Sb I \ As Na3AS I \ Te Na3P I \ s Na3N I \ 12 ~c1F-OF2-NF3 -CCl4-9F3-8eF2- Na20 I.. \ F2 - IF1 - SFs - PFs - SiF4 -AIF 3 - Mg!=2 - CsF Covalent Ionic Figure 2.1 Triangle illustrating the transitions between ionic, covalent and metallic bonding. (Reproduced from Chemical Constitution, by J.A.A. Ketelaar, Elsevier.) l]I]c===~~~~~~~-IN_T_R_O_D_U_C_T_l_O_N_T_O_B_O_N_D_I_N_G~~ TRANSITIONS BETWEEN THE MAIN TYPES OF BONDING Few bonds are purely ionic, covalent or metallic. Most are intermediate between the three main types, and show some properties of at least two, and sometimes of all three types. · Ionic bonds Ionic bonds are formed when electropositive elements react with electro- negative elements. Consider the ionic compound sodium chloride. A sodium atom has the electronic configuration ls 2 2s 2 2p 6 3s 1 The first and second shells of electrons are full, but the third shell contains only one electron. When this atom reacts it will do so in such a way that it attains a stable electron configuration. The noble gases have a stable electron arrangement and the nearest noble gas to sodium is neon, whose configuration is ls 2 2s2 2p6 If the sodium atom can lose one electron from its outer shell, it will attain this configuration and in doing so the sodilim acquires a net charge of + 1 and is called a sodium ion Na+. The positive charge arises because the nucleus contains 11 protons, each with a positive charge, but there are now only 10 electrons.. Sodium atoms tend to lose an electron in this way when they are is supplied with energy, and so sodium an ekctropositive element: Na - Na+ + electron sOuild up molecules in this way in order to understand their electronic str~ctures, it does not follow that the atoms will react together directly. In this case, carbon and chlorine do not react directly, and tetrachloromethane is made by indirect reactions. A molecule of ammonia NH 3 is made up of one nitrogen and three hydrogen atoms: N · + 3[H ·] -+ H N H H The nitrogen atom is three electrons short of a noble gas structure, and the hydrogen atoms are one electron short of a noble gas structure. Nitrogen forms three bonds, and the hydrogen atoms one bond each, so all four atoms attain a stable configuration. One pair of electrons on the N atom is not involved in bond formation, and this is called a lone pair of electrons. Other examples of covalent bonds include water (with two covalent bonds and two lone pairs of electrons), and hydrogen fluoride (one covalent bond and three lone pairs): H : O H F: H Oxidation numbers The oxidation nu~ber of an e.lement in a covalent compound is calculated by assigning shared electrons to the more. electronegative element, and then counting the theoreti.cal charge left c;m each atom. (Electronegativity is described in Chapter 6.) An alternative approach is to break up (theoretically) the molecule by removing all the atoms as ions, and counting the charge left on the central at.om. It must be emphasized that molecules are not really broken, nor electrons really moved. For example, in H 20, removal of two H+ leaves a charge of -2 on the oxygen atom, so the oxidation state of 0 in H 2 0 is· (-11). Similarly in H 2S the oxidation state of S is ( - II); in F20 the oxidation state of 0 is (+II); in SF4 the oxidation state of S is (+IV); whilst in SF6 the oxidation state of S is (+VI). The concept of oxidation numbers works equally well with ionic compounds. and in CrCl 3 the·cr atom has an oxidation state of (+III) and it forms Cr3 + ions. Similarly in CrCfi, Cr has the oxidation state (+II) , and exists as Cr2 + ions. ·. Coordinate bonds · a A covalent bond results from the sharing of ·paif of electrons between two atoms, where each atom contributes.,one electron to the bond. It is also  TRANSITIONS BETWEEN THE MAIN TYPES OF BONDING possible to have an electron pair bond where both electrons originate from one atom and none from the other. Such bonds are called coordinate bonds or dative bonds. Since, in coordinate bonds, two electrons are shared by two atoms, they differ from normal covalent bonds only in the way they are formed, and once formed they are identical to normal covalent· bonds. Even though the ammonia molecule has a stable electron configuration, it can react with a hydrogen ion H+ by donating a share in the.lone pair of electrons, fortn_ing the ammonium ion NHt: H H + H + I H N: + (Hj+ "-+ H N:H or H-N-+H I H H H Covalent bonds are usually shown as straight lines joining the two atoms, and coordinate bonds as arrows indicating which atom is donating the electrons. Similarly arnmonia may donate its lone pair to boron trifluoride, and by this means the boron atom attains a share in eight electrons: H F H F I I H N:+B F ·- H-N-+B-F I I H F H F In a similar way, a molecule of BF3 can form.a coordinate bond by accepting a share jfi a lone pair from a p- ion. There are many other examples, including: PCls + c1- - [PC16r SbFs + p- - [SbF6r Double and triple bonds Sometimes more than two electrons are shared between a pair of atoms. If tour electrons are shared, then there are two bottds, and this arrangement is called a double bond. If six electrons are shared then there are three bonds, and this ·is called a triple bond: ~'~~~~~~~~~-IN_T_R_O_D_U_CT~IO_N_T_O~BO_N_D_I_N_G~~~~~~~~--' H H H H.\ I c c C=C Ethene molecule H I \ (double bond) H H H H:C C:H H-C=C-H Ethyne molecule (triple bond) MeraJlic bonds and metallic structures Metals are made up of positive ions packed together, usually in one of the three following arrangements: 1. Cubic close-packed (also called face-centred cubic). 2. Hexagonal close-paclced. _3. Body-centred cubic. Negatively charged electrons hold the ions together. The number of positive and negative charges are exactly balanced, as the electrons originated from the neutral metal atoms. The outstanding feature of metals is their extremely high electrical conductivity and thermal conductivity, both of which are because of the mobility of these electrons through the lattice. The arrangements of atoms in the three common metallic structures are shown in Figure 2.2. Two of these arrangements (cubic close-packed and hexagonal close-packed) are based OQ the closest packing of spheres. The metal ions are assumed to be spherical, and are packed together to fill the space most effectively, as shown in Figure 2.3a. Each sphere touches six other spheres within this one layer. A second layer of spheres is arranged on top of the first layer, the protruding parts of the second layer fitting into the hollows in the first layer as shown in Figure 2.4a. A sphere in the first layer touches three spheres in the layer above it, and similarly touches three spheres in the layer below it, plus six spheres in its own layer, making a total of 12. The coordination number, or -number of atoms or ions in contact with a given atom, is therefore 12 for a close-packed. arrangement. With a close-packed arrange- ment, the spheres occupy 74% of the total space. When adding a third layer of spheres, two different arrangements are possible, each preserving the close-packed arrangement. If the first sphere of the third layer is placed in the depression X shown in Figure 2.4a, then this sphere is exactly above a sphere in the first layer. It follows that every sphere in the third layer is exactly above asphere in the first layer as shown in Figure 2.2a. lfthe first layer is represented by A, and the second layer by H, the repeating pattern of close-packed sheets is AB AS AB ;... This structure has hexagonal symmetry, and it is therefore said to be hexagonal close-packed.. · Alternatively, the first sphere of the third layer may be placed in a  '--~~~~-T_RA_N_S_l_Tl_O_N_S_B_E_TW~E_EN~T_H_E_M_A_l_N_T_Y_P_ES_O~F_BO_N~D_IN_G~~~~____J"37/ A B A (a) 3-fold axis I 3-fold axis I ~ A c \ \ A \ (b) \ J/ I (c) Figure 2.2 The three metallic structures. (a) Htxagonal close-packed structure showing the repeat pattern of layers ABABAB... and the 12 neighbours sur- rounding each sphere. (b) Cubic close-packed structure (cbordinatiott number is also 12) showing repeat pattern of layers ABCABC. (c) Body-centred cubit structure showing the 8 tteighbOuts surrounding each sphere. ·  0] l..__________l_N_T_RO_D_UCT_l_O_N_T_O_BO_N_D_l_N_G_ _ _ _ _ _ _ _ j (a) (b) Figure 2.3 Possible ways of packing equal spheres in two dimensions. (a) Close- packed (fills 74% of space). (b) Body-centred cubic (fills 68% of space). A B A (a) Figure 2.4 Superill)pp~d layers of d~·Pl1Cked spheres. (a) Two layers of close- packed spheres (second layer is shaded). (b) Three layers of close-packed spheres (second layer shaded, third layer bold circles). Note that the third layer is not above the first layer. hence this is an ABCABC... (cubic close-packed) arrangement. A B depression such as Y in Figure 2.4a. The sphere is not exactly above a sphere in the first layer, and it follows that all the spheres in the third layer are not exactly above spheres in the first layer (Figure 2.4b). If the three c layers are represented by A. Band C. then the repeating pattern of sheets is ABCABCABC... (Figure 2.2b). This structure has cubic symmetry and (b) is said to be cubic close-packed. An. alternative name for this structure is face-centred cubic. The difference between hexagonal and cubic close :.5 Arrangement of 12 packing is illustrated in Figure 2.5. neighbours in hexagonal ic close-packed Random forms of close packing such as ABABC or ACBACB are ments. (Note that the top possible. but occur only rarely. Hexagonal ABABAB and cubic ABCABC die layers are the same. close packing are common. · 1e cubic close-packed The third common metallic structure is called body-c~ntred cubic e the bottom layer is (Figure 2.2c). The spheres are packed in sheets as shown in Figure 2.3b. 50° relative to the 1al close-packed. (a) The second layer occupies the hollows in this first sheet. The third layer nal close-packed. {b) occupies hollows in the second layer. and the third layer is immediately ose-packed. above the first layer. This form of packing is less efficient at filling the s,pace  '-------T_R_A_N_Sl_T_IO_N_S_B_E_TW_E_E_N_,_. T_H_E_M_A_IN_T_YP_E_S_O_F_B_O_N_D_l_N_G_ _ _ _ ___,j j 39 j than closest packing (compare Figures 2.3a and b). In a body-centred cubic structure the spheres occupy 68% of the total space and have a co- ordination number of 8, compared, with close-packed structures where 74% of the space is occupied and the coordination number is 12. Metallic structures always have high coordination numbers. The theories of bondjng in metals and alloys are described in Chapter 5. Metallic bonding is found not only in metals and alloys, but also in several other types of compound: 1. Interstitial borides, carbides, nitrides and hydrides formed by the transition elements (and by some of the lanthanides too). Some low oxidation states of transition metal halides also belong to this group, where the compounds show electricalconductivity, and are thought to contain free electrons in conduction bands. 2. Metal cluster compounds of the transition metals, and cluster com- pounds of boron, where the covalent bonding is delocalized over several atoms, and is equivalent to a restricted fortn of metallic bonding. 3. A group of compounds including the.metal Carbonyls which contain a metal-metal bond. The cluster compounds, and the compounds with metal-metal bonds, may help to explain the role of metals as catalysts. ,Melting points Ionic compounds are typically solids and usually have high melting and boiling points. In contrast covalent compounds are typically gases, liquids or low melting solids. These differences occut because of differences in bonding and structure. ·.... ·. · Ionic compounds,are made up of positive and negative ions arranged in a regular way in a lattice. The attraction between ions is electrostatic, and is non-directional, extending equally fo all directio11s. Melting the compound involves breaking the lattiee. This requires cohsiderable energy, and so the melting point and boiling point are usually high, and the compounds are very hard. · Compounds with covalent bonds are usually made up of discrete molecules. The bonds are directional, and strong covalent bonding forces hold the atoms together to make a molecule. In the solid, molecules are held together by weak van der Waals forces. To melt or boil the compound we only need supply the small amount of en~rgy needed to break the van der Waals forces. Hence covalently bonded compounds are often gases, liquids or soft solids with low melting points. In a few cases such as diamond, or silica Si02 , the structures are covalent giant lattices instead of discrete molecules. In these cases there is a three- dimensional lattice, with strong covalent bonds in all directions: It requires a large amount of energy to break.this lattice, arid so _diamond; silica and other materials with giant three-dimensional lattices are very hard and have high meltifig points. GI] j~_________IN_T_R_O_D_UCTION TO BONDIN~---~~===~-~~-~: ::-_ :J - · -·.. __ ___ Conductivity Ionic compounds conduct electricity when the compound is melted, or in solution. Conduction is achieved by the ions migrating towards the electrodes under the influence of an electric potential. If an electric current is passed through a solution of sodium chloride, Na+ ion_s are attracted to the negatively charged electrode (cathode), where they gain an electron and form sodium atoms. The c1- ions are attracted to the positive electrode (anode), where they lose an electron and become chlorine atoms. This process is called electrolysis. The changes amount to the transfer of electrons from cathode to unodc, but conduction occurs hy an ionic mechanism involving the migration of both positive and negative ions in opposite directions. In the solid state, the ions are trapped in fixed places in the crystal lattice, and as they cannot migrate, they cannot conduct electricity in this way. It is, however, wrong to say that ionic solids do not conduct electricity without qualifying the statement. The crystal may conduct electricity to a very small extent by scmiconduction if the crystal contains some defects. Suppose that a lattice site is unoccupic+ 0- Lone pairs occupy two Of the corners, and F atoms occupy the other three corners. Three different arrangements are theoretically possible, as shown in Fig1,1re 4.4. · The most stable structure will be the one of lowest energy, that is the one with the minimum repulsion between the five orbitals. The greatest repulsion occurs between two lone pairs. Lone pair-bond pair repulsions are next strongest, and bond pair-bond pair repulsions the weakest... F (1)... (2) Figure 4.4 Chlorine trifluoride molecule.  l~~~~~~~~_s_o_M_E~EX_A_M~P_LE_S~U_SI_N_G~T_H_E_V_S_E_P_R_T_H_E_O_R_Y~~~~~~~ ~ Groups at 90° to each other repel each other strongly, whilst groups 120° apart repel each other much less. Structure 1 is the most symmetrical, but has six 90° repulsions between lone pairs and and atoms. Structure 2 has one 90° repulsion between two lone pairs, plus thre~ 90° repulsions between lone pairs and atoms. Struc- ture 3 has four 90° repulsions between lone pairs and at