BONDING AND STRUCTURE

Transcription

1 BONDING AND STRUCTURE We are living in a world with millions of substances having properties vastly different from one another. It is the way the atoms are held together (bonding) and how they are arranged in space (structure) that accounts for the quite distinct properties of these different substances. When elements form compounds, they either lose, gain or share electrons so as to achieve stable electron configurations. This simple idea forms the basis of the electronic theory of bonding : a chemical bond is a special arrangement of electrons between atoms by which the resulting nuclei and electrons become more stable. There are three main types of bonding ionic, covalent and metallic which involves respectively transfer, sharing and pooling of electrons between atoms to achieve stable electronic arrangements. Type of bonds IONIC BOND COVALENT BOND METALLIC BOND formed between metals and nonmetals way to achieve stable electronic arrangement nondirectional nature of bonds electrostatic attraction between nuclei and shared electrons structure adopted example sodium chloride iodine diamond sodium nature of forces holding particles together strong metallic bond between ions and sea of electrons properties : (1) m.p. / b.p. (2) electrical conductivity (3) solubility IONIC BONDING A simple model of ionic bonding where electrons are being transferred from metal atoms to nonmetal atoms so that the resulting ions obtained a full outer shell of electrons :

2 Ionic Radii Bonding 2 The electron cloud of any atom or ion has no definite limit, thus the size of an atom cannot be defined in a simple and unique manner. We could, however, measure very precisely the distance between two nuclei by means of electron density maps. The figure shows the electron density map for sodium chloride. (a) Decide which is the sodium ion and which is chloride ion. (b) Use a ruler and the map scale to obtain approximate values for the ionic radii of sodium and chloride ions. The sizes of ions can be conveniently compared by measuring their ionic radii from electron density maps. 1 Comparison of sizes of ions with their parent atoms : (i) cations are smaller than their parent atoms (removal of a complete outer shell leads to a contraction of electron cloud) (ii) anions are larger than their parent atoms (affinity of incoming electrons results in an expansion of electron cloud) cations are usually smaller than anions 2 Comparison of sizes of isoelectronic particles : In any isoelectronic series, all the particles carry the same number of electrons but have a progressively increasing number of protons. Along an isoelectronic series the p/e ratio (no. of protons / no. of electrons) increase effective nuclear charge increase result in a contraction of the electron cloud of the ion ionic radii decrease along an isoelectronic series Structure of Ionic Lattices There are two common arrangements for most ionic compounds, namely the sodium chloride and caesium chloride structures. They can be conveniently represented by unit cell, defined as the smallest portion of structure which when repeatedly stacked together at all directions, can reproduce the entire lattice. It is made up of two interpenetrating facecentered cubic lattices (f.c.c.) of each type of ion (Na, Cl ) to form a simple cubic lattice with each type of ion occupying alternate corners It is made up of two interpenetrating simple lattices of each type of ion (Cs, Cl ) to form a bodycentered cubic lattice (b.c.c.) NaCl structure CsCl structure (a) Sketch the unit cell of the two different structures.

3 (b) Deduce the number of each type of ion per unit cell for each structure. NaCl structure : CsCl structure : Bonding 3 (c) State the respective coordination numbers of each type of ion in the two structures. NaCl structure : CsCl structure : Theoretically, ionic structures tend to have as higher coordination numbers as possible, because it allows maximum attractive forces operate between oppositely charged ions, making the resulting lattices more stable. Sodium chloride, however, does not adopt the caesium chloride structure because of the difference in relative ionic sizes. However, Na ions, being much smaller than Cs ions, cannot accommodate more than 6 Cl ions before these anions repel each other too strongly for a stable arrangement. Study the following two unit cells and deduce which one belong to calcium fluoride and which one to zinc sulphide. Physical Properties of Ionic Compounds Compounds of ionic solids are : 1 nonvolatile with high melting points and high boiling points 2 good conductors of electricity when molten, but nonconductors when solid 3 soluble in water, but insoluble in nonaqueous solvents such as tetrachloromethane (CCl 4 ) 4 hard and brittle Ionic crystals are hard because of the strong attractions between opposite charged ions. However, they are also very brittle and may be split clearly (cleaved) using a sharpedged razor. When the crystal is tapped sharply along a particular plane it is possible to displace one layer of ions relative to the next. As a result of this displacement, ions of similar charge come together. Repulsion then occurs, forcing apart the two portions of the crystal.

4 COVALENT BONDING According to valence bond theory, a covalent bond is formed by sharing the valence (outermost) electrons : Bonding 4 This model is inadequate because it suggests a pair of electrons to be fixed between the two nuclei, yet it is impossible to locate the exact position of electrons. A better model using the idea of overlapping orbitals/charge clouds is adopted : The actual charge distribution for a covalent molecule could be obtained from an electron density map, which shows a region of negative charge (shared electrons) located between two positive charges (nuclei). The covalent radius of an atom is defined as half the internuclear distance between two covalently bonded atoms in the molecule of the element. What is the covalent radius of hydrogen? Despite the limitation of valence bond theory, it is nevertheless very useful in making predictions about the electronic structures of covalent molecules by dotandcross diagrams. Draw electronic structures for the following species : (a) NCl 3 (b) H 2 S (c) CHCl 3 (d) C 2 H 4 (e) N 2 (f) CO 2 For polyatomic ions, such as NH 4 and CO 3 2, it is important to distinguish between the ionic bonding which binds these ions to other ions, and the covalent bonding which binds the atoms within each ion. Draw electronic structures for the following species : (a) OH (b) CO 3 2 (c) HCO 3

5 Dative Covalent Bond Bonding 5 In a normal covalent bond, each atom provides one electron for the shared pair. In a few compounds, however, a bond is formed by the sharing of an electron pair which are provided by one atom. This kind of bonding is known as dative covalent bond or coordinate bond, since both electrons in the bond are donated by one atom. Ordinary covalent bonds are represented by a short line as a shorthand for an electron pair. Dative covalent bonds are represented by an arrow showing the direction of donation. However, once a dative bond is formed, you should not imagine them being different from ordinary bonds, because electrons are identical irrespective of where they are coming from. Draw electronic structures for the following species : (a) CO (b) NH 4 (c) NO 3 Limitation of Octet Rule So far, we have considered only those covalent compounds which obey the octet rule. However, there are a number of cases where the rule does not apply. 1 Compounds with more than eight valence electrons per atom The octet rule always applies for elements in the first two periods of the Periodic Table : each atom cannot hold more than eight electrons in its valence shell. Starting from period 3, however, elements could make use of their vacant dorbitals to accommodate more than 8 valence electrons, apparently breaking the octet rule. Draw electronic structures for the following species : (a) PCl 5 (b) SO 4 2 (c) PO Compounds with fewer than eight valence electrons per atom When elements with fewer than four outer shell electrons form compounds they usually lose those electrons to form ions. For small atoms, however, the relevant ionization energies may be so high that covalent bonding occurs instead. Since there are fewer than four electrons available for sharing, there will then be fewer than eight outer shell electrons per atom in the resulting compound. Such compounds are often called electrondeficient. Draw electronic structures for the following species : (a) BF 3 (b) BeCl 2 (c) AlCl 3 There is a tendency for these electrondeficient compounds to achieve 'stable octet'by accepting a lone pair of electrons from neighbouring molecules to form a dative covalent bond. NH 3 BF 3 BeCl 2 (polymer) Al 2 Cl 6 (dimer)

6 Shapes of Molecules The electronpair repulsion theory allows us to make predictions about the shapes of a great many covalent molecules : Bonding 6 1 Each covalent bond is regarded as a pair of electrons. An electron pair may be a shared pair, lone pair, single bond, double bond or triple bond. 2 Electron pairs around a central atom tend to get as far apart from each other as possible, so as to minimize the electrostatic repulsion among them. 3 Nonbonding pair of electrons exerts a greater repelling effect than bonding pair of electrons does. Therefore, lone pair lone pair lone pair bond pair bond pair bond pair repulsion repulsion repulsion no. of electron pairs around central atom 2 Basic skeleton Bond angle examples BeCl 2 CO 2 linear BF 3 SO 3 SO 2 3 trigonal planar CH 4 NH 3 H 2 O 4 tetrahedral PCl 5 ClF 3 ICl 2 5 trigonal bipyramidal SF 6 BrF 5 XeF 4 6 octahedral

7 Bonding 7 For each of the following molecules, draw a threedimensional structure showing the positions of the bond electron pairs and lone electron pairs (if any) of the central atom. (i) PCl 4 (ii) PCl 6 (iii) PCl 3 A few comments on bond angles : 1 CH 4 molecules are perfectly symmetrical. The bond angle for HCH is that of a regular tetrahedron, Although sharing the same molecular skeleton, NH 3 and H 2 O have different bond angles (107 for HNH and 105 for HOH). The region in space occupied by a nonbonding pair of electrons is smaller and closer to the nucleus of an atom than a bonding pair. Bonded pairs of electrons are drawn out between the nuclei of the two covalentlybonded atoms. This means that a lone pair can exert a greater repelling effect than a bonded pair and this results in a decreasing bond angle from CH 4 to NH 3 to H 2 O. 2 NH 3 and PH 3 are both pyramidal with four pairs of electrons around the central atom, one of which is a lonepair. The bond angle for HNH in NH 3 is 107 while that for HPH is 90. N atom is smaller than P. It follows that NH bond pairs lie much closer to the central atom than PH bond pairs do. The repulsion among NH bond pairs are thus much greater than that among PH bond pairs, resulting a decrease in bond angle down the group. Giant Covalent Structures The sharing of electron pairs does not necessarily result in discrete molecules. In a few cases, covalent bonding is extended indefinitely in three dimensions. Such giant structure is well illustrated by the two allotropes (different structural forms of the same element) of carbon, namely diamond and graphite. The properties of diamond and graphite are in dramatic contrast. Since both solids consists of identical carbon atoms only, the differences in properties must be entirely due to differences in structures. DIAMOND GRAPHITE structure coordination no. / hybridization m.p. / b.p. Each C atom is covalently bonded to 4 other C atoms in a tetrahedral manner to form a 3D giant network Each C atom is covalently bonded to 3 other C atoms in trigonal planar to form a multilayer structure. Adjacent layers are held by van der Waals'forces 4 / sp 3 3 / sp C 3700 C strong CC covalent bonds have to be broken in melting / boiling strength hard; strong & directional CC bonds restrict relative motion between C atoms soft; weak van der Waals' forces allow layers to slip over each other easily electrical conductivity nonconductor of electricity; all electrons are localized in covalent bonds good conductor of electricity; each C atom has a delocalized electron which can move freely along the same layer uses jewelries, glass cutters electrodes, lubricants, pencils

8 INTERMEDIATE TYPE OF BONDING In many compounds, the bonding is intermediate in character between 'pure' ionic and 'pure'covalent. We can discuss intermediate type of bonding in two ways, either by seeing : 1 how ordinary ionic bonds can have covalent character; and 2 how ordinary covalent bonds can have ionic character. Bonding 8 Covalent Character of Ionic Bonds The model of ionic bonding so far assumes the complete transfer of electrons from one atom to another giving separate spherical ions. However, the spherical charge cloud in large anions may be easily distorted, or polarized, by neighbouring small cations, because the nuclear charge of the anion is insufficient to hold its electron cloud firmly. Such distortion leads to a certain overlap of charge clouds between cations and anions, and hence a degree of covalent character. The distortion of an electron cloud by a neighbouring charged particle is called polarization. Evidence of whether an ionic compound has significant covalent character comes from a comparison of the lattice energy values, one derived from theory and the other determined from experiment. Assumptions in deriving the theoretical lattice energy of an ionic compound : 1 ions are spherical 2 charge is uniformly distributed throughout the ions 3 forces operating between ions are only electrostatic in nature theoretical value applying Law of Electrostatics experimental value applying Hess'Law underlying principle L.E. q1 q2 r L.E. = B C D E A NaCl LiF AgCl ZnS Cations such as Al 3, Ag and Zn 2 have small ionic radii and high nuclear charge, both give rise to a high charge density (i.e. charge/volume ratio), thus distorting the electron cloud of neighbouring anions to a greater extent. These ions are said to possess a high polarizing power. Electrons in larger anions are further away from the nucleus and are less firmly held by the nucleus than smaller anions. Large anions are said to have a high polarizability (the ability to become polarized).

9 Ionic Character of Covalent Bonds Bonding 9 Polar covalent bond and dipole moment The model of covalent bonding used so far assumes the equal sharing of a pair of electrons between two atoms. Very often the electrons are not shared equally, as in the case when a bonding pair of electrons is attracted more to one nucleus than the other (e.g. Cl atom in HCl molecule is more electronegative than the H atom). This makes the centre of negative charge not coincide with the centre of positive charge, resulting in the formation of a dipole with two equal and opposite charges (q and q) separated by a distance d This bond is said to be polarized and described as a polar covalent bond. There are four ways of representing bond polarity : The extent of bond polarization can be measured in terms of dipole moment, µ, which is given by µ = q x d. Dipole moments are usually expressed in Debye (D) unit (e.g. dipole moment of HCl molecule is 1.1 D) 1 Dipole moments are vector quantities. For a molecule with more than one polar bond, the dipole moment is given by the vector sum of the dipole moments of various polar bonds. If the vector sum is zero, the dipole moment of the molecule is zero, and the molecule is described as nonpolar. 2 The greater the overall dipole moment, the more polar the molecule is. 3 Polarity of a liquid can be tested by studying the effect of a charged rod on a stream of liquid from a burette. Any deflection of the stream indicates that the liquid is polar. 4 Dipole moments can provide useful information about the structure of molecules. As an example, the zero dipole moment of CO 2 shows that the molecule must be linear such that the dipole moment of each C=O bond cancels each other out. On the other hand, the existence of a net dipole moment for SO 2 molecule indicates that the molecule contains polar bonds which are not linearly arranged. By drawing a 3D structure showing the bond electron pairs and lone electron pairs (if any), state whether or not the molecule possesses a dipole moment. (i) BCl 3 (b.p.13 C) (iii) CCl 4 (b.p.77 C) (v) CO 2 (b.p.78 C) (ii) NCl 3 (b.p.71 C) (iv) CHCl 3 (b.p.62 C) (vi) SO 2 (b.p.10 C) Electronegativity In polar covalent bonds, the unequal sharing of the bonded electron pairs is caused by a difference in electron attracting ability of the bonded atoms. This electron attracting ability of the atom is described as electronegativity of the element. 1 The Pauling's scale is used to measure electronegativities of different elements : (i) a value of 4.0 is assigned to the most electronegative element, fluorine (ii) all other elements having a lower electron attracting ability than fluorine are assigned lower electronegativity values

10 Bonding 10 2 Smaller atoms with greater charge density tend to attract electrons more tightly than larger ones, i.e. electronegativity increases across a period and decreases down the group 3 Electronegativity differences govern the nature of chemical bonding bond electronegativity difference between bonded atoms type of bond Cl Cl = 0 covalent Cl H = 0.9 polar covalent F Li = 3.0 ionic 4 Both electronegativity and electron affinity refer to the attraction by an atom for electrons. The essential difference between these terms lies in the fact that electron affinity refers to the attraction between incoming electrons (external) and isolated atoms; whereas electronegativity refers to the attraction between bonding electrons (internal) and an atom in a covalent compound. INTERMOLECULAR FORCES There are three major types of interactions operating between simple covalent molecules, depending on the extent of polarity of the molecules : 1 permanent dipoledipole interactions 2 dipoleinduced dipole interactions 3 temporary (induced) dipoleinduced dipole interactions (or van der Waals forces in general) Dipole Dipole Interaction Dipole Induced Dipole Interaction Polar molecules, or sometimes called dipoles, are attracted to one another when the positive end of one molecule is oriented toward the negative end of its neighbour. Such electrostatic attractions between permanent dipoles are considerably strong, and explains why polar compounds usually have relatively higher melting or boiling points than nonpolar molecules of similar molecular size. It arises when a polar molecule, in the vicinity of another molecule (which is nonpolar), has the effect of polarizing the second molecule. The induced dipole can then interact with the first molecule. Induced Dipole Induced Dipole Interaction Nonpolar molecules, though having no permanent dipole moment, do have temporary, constant changing dipoles due to the continuous motion of the electrons. 1 Electrostatic attractions between temporary induced dipoles are weaker than that between permanent dipoles. 2 Electron cloud of larger molecules are more readily to be polarized than smaller ones, resulting in a stronger intermolecular attraction. This is why substances with larger molecular size generally have higher melting and boiling points. 3 Temporary induced dipoles are present in BOTH polar and nonpolar molecules (as long as they have polarizable electron clouds). For large polar molecules, the attractive forces due to temporary induced dipoles may even be greater than those due to the permanent dipoles.

11 Bonding 11 Hydrogen Bonding Hydrogen bonding is often regarded as a special kind of permanent dipoledipole attractions. This is formed between a lone pair of electrons from a highly electronegative element (ie. F, O or N) and a hydrogen atom, which itself bonded to another very electronegative element. δ δ δ δ δ δ F H F O H N N H O Hydrogen bond is by far the strongest type of intermolecular forces, with strength about onetenth of an ordinary covalent bond. This relative strength has some important and interesting consequences : 1 Boiling points of hydrides of group V, VI and VII In general, boiling point increases down the group, owing to the increase in molecular size ( greater v.d.w. forces) However, the first member of the hydride from Group V, VI & VII have exceptionally high boiling points, because of the existence of intermolecular hydrogen bonds Such abnormal phenomenon does not appear in Group IV hydrides, as the first member here (i.e. carbon) is obviously not electronegative enough to form intermolecular Hbonds 2 Structure and density of Ice Abnormal behaviour of ice : 1 Ice floats on water while solid HF does not float on its own liquid. 2 When liquid water is heated from 0 C, its density increases for a short while (below 4 C) before decreasing in the usual way. Explanation : H 2 O has two lone pairs of electrons and could form at maximum two hydrogen bonds per molecule. Each H 2 O molecule is, therefore, surrounded tetrahedrally (2 covalent bonds and 2 hydrogen bonds) to form an 'open' cagelike structure. As ice melts, some of these hydrogen bonds are broken. The partial collapse of such an open structure allows water molecules to pack closer to each other density of water > density of ice 3 Hydrogen bondings in proteins and DNA Hydrogen bonding plays a crucial role in the structures of many biochemical substances such as proteins and DNA : Protein chains are held in close proximity to one another by hydrogen bonds formed between the NH and C=O groups of the neighbouring chains In DNA, the two helical strands of nucleic acid chains, which controls reproduction and inheritance, are linked together by a series of hydrogen bonds

12 external pressure 4 Relative viscosity Viscosity of a liquid is related to the ease with which molecules can move past one another in the liquid. Viscosity of alcohols increases in the following order : propan1ol < propane1,2diol < propane1,2,3triol Bonding 12 Each alcohol molecule forms hydrogen bonds with its neighbours through hydroxyl (OH) groups. As the number of hydroxyl groups increases, so does the extent of hydrogen bonding in the liquid; making it more viscous. 5 Pressuretemperature diagram vapour pressure of a liquid at a particular temp. is the pressure exerted by the vapour molecules, which are in equilibrium with the liquid molecules at that temp. graphical representation : at each temp., the vapour pressure of the more volatile substance is always higher than the less volatile one boiling point is defined as the temp. at which the vapour pressure of the substance is equal to the atmospheric pressure note that liquids with lower vapour pressures (ie. less volatile) will have correspondingly higher b.p. vapour pressure temperature the curve represents the variation of boiling points with different pressure and temp. (notice that the yaxis now indicates external pressure rather than vapour pressure of the substance) each area included by the curve represents the conditions of temp. and pressure under which a particular phase is stable (or dominant) temperature every point on the curve represents the conditions of temp. & pressure at which the two phases coexist in equilibrium the concept can be further elaborated to represent equilibria involving all the 3 phases : interpretation of phase diagram AB : BC : BD : B : point (at which all the three phases coexist in equilibrium) C : point (above which the vapour cannot be liquefied no matter how high the applied pressure will be) Determine each of the following by referring to the above phase diagram of a substance, X : (a) the m.p. and b.p. of X at atmospheric pressure ; (b) the state of X at (I) 200 C and 2 atm; (ii) 0 C and 1 atm ; (c) the condition under which X will sublime.

13 comparison of phase diagrams between CO 2 and H 2 O Bonding 13 P P features : consequences : T (i) triple point for CO 2 is above 1 atm while that for water is below 1 atm (ii) slope of solidliquid curve for CO 2 is ve while that for water is ve (i) liquid CO 2 is not stable under ordinary conditions solid carbon dioxide sublimes rather than melts (ii) an increase in P will lower m.p. for H 2 O delay freezing or melting occurs earlier than before water freezes with expansion in volume or ice melts with contraction in volume water is denser than ice (formation of open cage structure in ice) T Molecular Crystals Molecular crystals consist of molecules held in simple cubic or facecentred cubic lattice by weak intermolecular forces such as van der Waals' forces (e.g. iodine) or hydrogen bonds (e.g. ice) : Molecular solids are : 1 volatile with low m.p. / b.p. 2 poor conductor of electricity 3 usually soft and of low densities Buckministerfullerene, with formula C 60, was discovered in 1985 as the third form of carbon allotrope, which is made from interlocking hexagonal and pentagonal rings of C atoms. It is the parent of a new family of structures called fullerenes, including C 70, C 82 and C 100, mostly produced from the sooty flames from burning benzene. Fullerenes can conduct electricity due to the presence of delocalized electrons on their surface. By housing various metal atoms (e.g. K), inside the hollow cage, a whole new range of applications have been found, including catalysts, superconductors, rocket fuels, laser materials and drugs. Deduce the number of carbon atoms in a unit cell of C 60, if it is known to have a facecentred cubic structure.

14 METALLIC BONDING Bonding 14 The structure of a metal can be conveniently illustrated by a model in which a lattice of regularly packed cations is surrounded by a sea of delocalized electrons. 1 The delocalized nature of electrons accounts for their high electrical and thermal conductivity, whereas the flexibility of packing the cationic lattice accounts for their malleable (pressed easily) and ductile (pulled easily) behaviour 2 Strength of metallic bonding is governed by two factors : (a) no. of valence electrons metallic bond strength increases with the number of outermost (valence) electrons m.p./b.p. increases across the period (i.e. Na < Mg < Al) (b) metallic radii ionic radius increases down the group, resulting in a decrease in charge density and, therefore, a reduction of electrostatic attraction between cations and the valence electrons m.p./b.p. decreases down the group (i.e. Li > Na > K) 3 When a metal is melted, only a small portion of metallic bonds are broken, resulting in a slight distortion of the entire structure. On the other hand, all metallic bonds are broken in boiling, making the melting points of metals are usually significantly lower than their boiling points. (e.g. m.p. & b.p. for Na are 98 and 883 C respectively) Metallic Crystals The problem of packing cation together in a lattice is not so very different from the problem of packing pingpong balls into a box. There are two different ways of packing spheres within a single layer : close packing loose packing 1 In building up metal spheres in a 3D closepacked structure : a b In a closepacked layer, each atom is in contact with six others, resulting a hexagonal arrangement of atoms. In the second layer, atoms pack as closely as possible to those in the first layer by 'sitting' in the holes between atoms in the first layer. (Notice that there are now two different kinds of holes, namely tetrahedral and octahedral, depending on the shape formed by the neighbouring spheres surrounding the hole) c The third layer of atoms can now be added in two quite distinct ways so that : (i) each sphere in the layer is sitting on the tetrahedral holes formed by the previous two layers, resulting in an overall pattern that repeats itself every alternate layer (denoted by abab... etc.) hexagonal close packing (h.c.p.) (ii) each sphere in the layer is sitting on the octahedral holes formed by the previous two layers, resulting in an overall pattern that repeats repeats itself every three layers (denoted by abcabc... etc.) cubic close packing (c.c.p.) (the word cubic is derived from the fact that a c.c.p. pattern could be visualized as a facecentred cubic (f.c.c.) structure from a different point of view)

15 Bonding 15 2 In building up metal spheres in a threedimensional open (not closepacked) manner, a second layer can be formed by placing a layer of spheres in the holes formed by the first layer, and a third layer formed in a similar way. This produces a bodycentred cubic (b.c.c.) arrangement in which each sphere is in contact with 4 spheres in the layer above and 4 in the layer below, but in contact with no spheres within its own layer. Different metals adopt different types of structures depending on their metallic radii. h.c.p. c.c.p. / f.c.c. b.c.c. arrangement for atoms in different layers ABAB ABCABC unit cell representation coordination number packing efficiency closely packed loosely packed Alloys An alloy is a mixture of two or more metals (or of a metal with a nonmetal) such that the resulting properties are generally more desirable than the pure metals alone. While some alloys (e.g. PbSn alloy solder) exist in heterogeneous mixtures with separate solid phases, most alloys are homogeneous solid solutions and can be further classified into: 1 substitutional alloy: some of the metal atoms being substituted by another metal (e.g. 18carat gold alloy with silver); atomic radii of the metals in the alloy should match with each other 2 interstitial alloy: the smaller atoms occupy holes among lattice formed by larger atoms (e.g. 1% carbon among iron in steel) Compared with their constituents, the properties of alloys can be modified in the following ways: Alloys are harder and stronger than pure metals. In the case of pure atoms, layers of atoms with identical radius can easily slide over each other. However, the addition of a small amount of one metal with different atomic radii disrupts the orderly arrangement of atoms, making it more difficult to slide over each other. (Degree of hardness increase as the relative motion between particles become more restricted.) some alloys are more resistant to corrosion than pure metals (e.g. stainless steel, AlMg alloy duralumin) many alloys form mixtures with much lower melting points than pure metals (e.g. solder to join metals together)