Inorganic Chemistry with Help Guides

  This cycle devised by Born and Haber. In 1919 relates the lattice energy of a crystal to other thermochemical data. The energy terms involved in building a crystal lattice such as sodium chloride may be taken in steps. The elements in their standard state are first converted to gaseous atoms, and then to ions, and finally packed into the crystal lattice.  The enthalpies of sublimation and dissociation and the ionization energy are positive since energy is supplied to the system. The electron affinity and lattice energy ate negative since energy is evolved in these processes. According to Hess's law, the overall energy change in a process depends only oil the energy of the initial and finl,ll states an,d not on the route taken. Thus the enthalpy of formation △Hr is equal to the sum of the terms going the other way round the cycle. 

  The energy released when an extra. electron is added to a neutrai gaseous atom is termed the electron affinity. Usually only one electron is added, forming a uninegative ion. This ~epels further electrons and energy is needed to add on a second electron: hence· the negative affinity of 0 2-. Electron affinities depend on the size and effective nuclear charge. They cannot be determined directly, but are obtained indirectly from the Born-Haber cycle.  Negative Clectron affinity values indicate that energy is.given out when the atom accepts an electron. The above values show that the halogens all evolve a large amount of energy on forming negative halide ions, and it is not surprising that these ions occur in. a large number of compounds. Energy is evolved when one electron is added to an 0 or S atom, forming the species o- and s-, but a substantial amount of energy is absorbed when two electrons are added to form 0 2- and s2- ions. Even though it requires energy to form these ...

  If a small amount of energy is supplied to an atom, then an electron may be promoted to a higher energy .level, but if the amount of. energy supplied is sufficiently large the electron may be completely removed. The energy required to remove the most loosely bound electron from an isolated gaseous atom is called the ionization energy. Ionization energies are determined from spectra and are measured in kJ mo1- 1 • It is possible to remove more than one electron from most atoms. The first ionization energy is the energy required to remove the first electron and convert M to M+; the second ionization energy is the energy required to remove the second election arid convert M+ to M2+; the third ionization energy converts M2+ to M3+, and so on. The factors that influence the ionization energy are:  1. The size of the atom.  2. The charge on the nucleus.  3. How effectively the inner electron shells screen the nuclear charge.  4. The type of electron involved (s. p, ...

 Irrespective of which set of ionic radii are used, the following trends are observed:  1. In the main groups, radii increase on descending the group, e.g. u+ = 0.76A, Na+ = l.02A, K+- = l.38A, because extra shells of electrons are added. 2. The ionic radii decrease moving from left to right across any period in the periodic table, e.g. Na+ = 1.02A, Mg2+ = 0.720Aand Al3+ = 0.535 A. This is partly due to the increased number of charges on the nucleus, and also to the increasing charge on the ions.  3. The ionic radius decreases as more electrons are ionized off, that is as the valency incteases, e.g. Cr2+ = 0.80 A (high spin), Cr3+ = 0.615 A, Cr4+ = 0.55 A, CrH = 0.49 A and Cr6+ = 0.44 A.  4. The d and f orbitals do not shield the nuclear charge very effectively. Thus there is a significant reduction in the size of ions just after 10d or 14f electrons have been filled in. The latter is called the lanthanide contraction, and results in the sizes of the second and third...

 There are several problems in obtaining an accurate set of ionic radii.  1. Though it is possible to measure the internuclear distances in a crystal very accurately by X-ray diffraction, for example the distance between Na+ and p- in NaF, there is no universally accepted formula for apportioning this to the two ions. Historically several different sets of ionic radii have been estimated. The main ones are by Goldschmidt. Pauling and Ahrens. These are all calculated from observed internuclear distances, but differ in the method used to split the distance between the ions. The most recent values, which are probably the most accurate, are by Shannon (1976) .  2. Corrections to these radii are necessary if the charge on the ion is changed.  3. Corrections must also be made for the coordination number, and the geometry.  4. The assumption that ions are spherical is probably true for ions from the s- and p-blocks with a noble gas configuration, but is probably untrue...

  Metals usually form positive ions. These are formed by removing one or more electrons from the metal atom. Metal ions are smaller than the atoms from which they were formed for two. reasons:  1. The whole of the ·outer shell of electrons is usually ionized, i.e. removed. This is one reason why catfons· are much smaller than the original metal atom.  2. A second factor is the effective nuclear charge. In an atom, the number of positive charges on the nucleus is exactly the same as the number of orbital electrons. When a positive ion is formed; the number of positive charges on the nucleus exceeds the number of orbital electrons, and the effective nuclear charge (which is the ratio of the nun:iber of charges on the nucleus to the number of electrons) is increased. This results in the remaining electrons being more strongly attracted by the nucleus. Thus the electrons are pulled in - further reducing the size.

  The size of atoms decreases fr()m left to right across a period in the periodic table. For example. on moving from lithium to beryllium one extra positive charge is added to the nucleus, and an extra orbital electron is also added. Increasing the nuclear c:harge results in all of the orbital electrons being pulled closer to the nucleus. In a given period, the alkali metal is the largest atom and the halogen the smallest. When a horizontal period contains ten transition elements the contraction in size is larger, and when in addition there are 14 inner transition eleme11ts in a horizontal period, the contraction in size is even more marked. On descending a group in the periodic table such as that containing lithium. sodium, potassiUm, rubidium and caesium, the sizes of the atoms increase due to the effect of extra shells of electrons being added: this outweighs the effect of increased nuclear charge.

 If two metals are completely miscible with each· other they can form a continuous range of solid solutions. Examples include Cu/Ni, Cu/Au, K/Rb, K/Cs and Rb/Cs. In cases like these, one atom may replace another at random in the lattice. In the Cu/Au case at teftlperatures above 450°C a disordered structure exists (Figure 5.7c), but on slow cooling the more ordered superlattice may be formed (Figure 5.7d). Only a few metals form this type Of continuous solid solution1 and Hume Rothery has shown that for complete miscibility the following three rules should apply.

  Next consider the relative sizes of the atoms. The structure of many metals is a close-packed lattice of spherical atoms or ions. There are therefore many tetrahedral and octahedral holes. If the element added has small atoms, they can be accommodated in these holes without altering the structure of the tnetal. Hydrogen is small enough to occupy tetrahedral holes, but most other elements occupy the larger octahedral holes. The invading atoms occupy interstitial.positions in the metal lattice, instead of replacing the metal atoms. The chemical composition of compounds Of this type may. vary over a wide range depending on how many holes are occupied. Such alloys are called interstitial solid solutions, and are formed by a: wide range of metals with hydrogen, boron, carbon, nitrogen and other eletrients'. The most li'tlportant factor is the size of the invading atoms. 

 When two metals are heatec;l together, or a metal is mixed with a nonmetallic element, then one of the following will occur:  l. An ionic compound may be formed.  2. An interstitial alloy may be formed.  3. A substitutional alloy may be formed.  4. A simple mixture may result.  Which of these occurs depends on the hemical nature of the two elements concerned, and on the relative sizes of the metal atoms and added atoms.

  In electrical conductors (metals), either the valence band is oniy partly fu1l, or the valence and conductiori bands overlap. There is therefore no significant gap between filled and un_filled MOs, and perturbation can occur readily. In insulators (non-metals), the valence band is full, so perturbation within the band is impossible, and there is an appreciable difference in energy (called the band gap) between the valence band and the next empty band. Electrons cannot therefore be promoted to an empty level where they could move freely. Intrinsic semiconductors are basically insulators, where the energy gap between adjacent bands is sufficiently small for thermal energy to be able to promote a small number of electrons from the full valence band to the empty conduction band. Both the promoted electron in the conduction band and the unpaired electron left in the valence band can conduct electricity. The conductivity of semiconductors increases with temperature, because the number ...

 Consider a simple m.etal such as lithium, which has a body-centred cubic structure,. with eight nearest neighbo11rs and six next-nearest neighbours at a slightly greater distance. A lithium atom has one electron in its outer shell, which may be shared with one of its neighbours, forming a normal two-electron bond. Th.e atom could eqQally well be bonded to any of its other eight neighbours, so many different arrangements are possible. A lithium atom may form two bonds if it ionizes, and it can then form many structures similar to those in ,Figures 5.lc and .d. Pauling suggested that the true structur~ is a mixture of all the many possible bonding forms. The more possible structures there are, the 19wer the energy. This means that the cohesive force which holds the structure together is large, and in metallic lithium the cohesive energy is three times greater than in a Li2 molecule. The cohesive energy increases from Group I to II to III, and this is explained by the atoms being abl...

  Inorganic chemistry   deals with synthesis and behavior of inorganic and organometallic compounds. This field covers chemical compounds that are not carbon-based, which are the subjects of organic chemistry. The distinction between the two disciplines is far from absolute, as there is much overlap in the subdiscipline of organometallic chemistry. It has applications in every aspect of the chemical industry, including catalysis, materials science, pigments, surfactants, coatings, medications, fuels, and agriculture. Key concepts Many inorganic compounds are ionic compounds, consisting of cations and anions joined by ionic bonding. Examples of salts (which are ionic compounds) are magnesium chloride MgCl2, which consists of magnesium cations Mg2+ and chloride anions Cl−; or sodium oxide Na2O, which consists of sodium cations Na+ and oxide anions O2−. In any salt, the proportions of the ions are such that the electric charges cancel out, so that the bulk compound is electricall...

  The honding and structures adopted hy metals and alloys arc less fully understood than those with ionic and covalent compounds. Any successful theory of metallic bonding must explain both the bonding between a large number of identical atoriJs· in a pure, metal, and the bonding between widely different metal atoms in alloys. The the.ory cahnot involve directional bonds, since most metallic properties remain even when the metal is in the liquid state (for example mercury), or when dissolved in a suitable solvent (for example solutions of sodium in liquid ammonia). Further, the theory should explain the great mobility of electrons.

  If the valence electrons in a metal are spread over a large number of bonds, each bond should be weaker and hence longer. The alkali metals exist as diatomic molecules in the vapour state, and the interatomic distances in the metal crystal are longer than in the diatomic molecule. Though the bonds in the metal are longer and weaker, there are many more of them than in the M2 molecule, so the total bonding energy is greater in the metal crystal. This can ·be seen by comparing the enthalpy of sublimation of the metal crystal with the enthalpy of dissociation of the M2 molecules.

  Metallic elements usually have a close-packed structure with a coordination number of 12. There are two types of close packing depending on the arrangement of adjacent layers in the structure: cubic close packing ABCABC and hexagonal close packing ABAB (see Metallic bonds and metallic structures in Chapter 2). However, some metals have a body centred cubic type of structure (which fills the space slightly less efficiently) where there are eight nearest neighbours, with another six next-nearest neighbours about 15% further away. If this small difference in distance between nearest and next-nearest neighbours is disregarded, the coor-dination number for a body-centred cubic structure may be regarded loosely as 14. The mechanical properties of malleability and ductility depend on the ease with which adj~nt planes of atoms can glide over each other, to give an equivalent arrangement of spheres. These properties are also affected by physical imperfections such as grain boundaries and ...

  Malleability and cohesive force   The mechanical properties of metals are that they are typically malleable and ductile. This snows that there is not much resistance to deformation of the structure. but that a large cohesive force holds the structure together. The cohesive force may be measured as the heat of atomization. Some numerical values of △H° , the heats Of atomization at 25°C, are given in Table. The heats of atomization (cohesive energy) decrease on descending a group in the periodic table Li-Na-K-Rb-Cs, showing that they are inversely proportional to the internudeat distatice. The cohesion energy increases across the periodic table from Group I to Group II to Group III. This suggests that the strefigth of metallic bonding is related to the number of valency electrons. the cohesive energy increases at first on crossing the transition series Sc-Ti-V as the number of unpaired d electrons increases. Continuing across the transition series the number of electrons per a...