5 Structure And Physical Properties 2009

STRUCTURE AND PHYSICAL PROPERTIES INTRODUCTION In the last chapter we have studied the covalent bond and the shape of small molecules with a reference to more complex structures. We can now predict the geometry of the particles in a covalent substance either molecular or giant covalent. In the case of ionic and metallic structure this geometric study is not relevant because the ionic and the metallic bond are essentially non-directional. Our next target is to study how the observable physical properties of pure substances depend on the interaction among its particles which in turn depends on their nature. We shall describe the arrangement of the particles in the different classes (ionic, metallic, molecular and macro-covalent) of materials and correlate some observable physical characteristics with the forces among these particles. IONIC STRUCTURES The structure of a typical ionic solid: sodium chloride How the ions are arranged in sodium chloride Sodium chloride is taken as a typical ionic compound although according to the charts of the previous section, it has a 67 % ionic character. Compounds like this consist of a giant (endlessly repeating) lattice of ions. So sodium chloride (and any other mostly ionic compound) is described as having a giant ionic structure. You should be clear that giant in this context doesn't just mean very large. It means that you can't state exactly how many ions there are. There could be billions of sodium ions and chloride ions packed together, or trillions, or whatever - it simply depends how big the crystal is. That is different from, say, a water molecule which always contains exactly 2 hydrogen atoms and one oxygen atom - never more and never less. A small representative bit of a sodium chloride lattice looks like this:

If you look at the diagram carefully, you will see that the sodium ions and chloride ions alternate with each other in each of the three dimensions. We normally draw an "exploded" version which looks like this: Only those ions joined by lines are actually “touching each other”. The sodium ion in the centre is being touched by 6 chloride ions. By chance we might just as well have centred the diagram on a chloride ion (that, of course, would be touched by 6

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sodium ions). Sodium chloride is described as being 6:6-co-ordinated. The pattern repeats in this way over countless ions. Why is sodium chloride 6:6-co-ordinated? To gain maximum stability, you need the maximum number of attractions. So why does each ion surround itself with 6 ions of the opposite charge? That represents the maximum number of chloride ions that you can fit around a central sodium ion before the chloride ions start” touching” each other. If they start touching, you introduce repulsions into the crystal which makes it less stable. In the case of a bigger central atom coordination can change. The physical properties of sodium chloride •

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Sodium chloride has a high melting and boiling point. There are strong electrostatic attractions between the positive and negative ions, and it takes a lot of heat energy to overcome them. Ionic substances all have high melting and boiling points. Differences between ionic substances will depend on things like: The number of charges on the ions: Magnesium oxide has exactly the same structure as sodium chloride, but a much higher melting and boiling point. The 2+ and 2- ions attract each other more strongly than 1+ attracts 1-. The sizes of the ions: If the ions are smaller they get closer together and so the electrostatic attractions are greater. Rubidium iodide, for example, melts and boils at slightly lower temperatures than sodium chloride, because both rubidium and iodide ions are bigger than sodium and chloride ions. The attractions are less between the bigger ions and so less heat energy is needed to separate them. Sodium chloride crystals are brittle: Brittleness is again typical of ionic substances. Imagine what happens to the crystal if a stress is applied which shifts the ion layers slightly. Ions of the same charge are brought side-by-side and so the crystal repels itself to pieces!

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Sodium chloride is soluble in water Many ionic solids are soluble in water although not all. It depends on whether there are big enough attractions between the water molecules and the ions to overcome the attractions between the ions themselves. Positive ions are attracted to the lone electron pairs on water molecules and co-ordinate (dative covalent) bonds may form. Water molecules form hydrogen bonds with negative ions (explained later). Sodium chloride is insoluble in organic solvents: This is also typical of ionic solids. The attractions between the solvent molecules and the ions aren't big enough to overcome the attractions holding the crystal together. The electrical behaviour of sodium chloride Solid sodium chloride doesn't conduct electricity, because there are no electrons which are free to move. Molten sodium chloride undergoes electrolysis, which involves conduction of electricity because of the movement of the ions. In the process, sodium and chlorine are produced. This is a chemical change rather than a physical process. 2

Water solutions of sodium chloride will conduct electricity for the same reasons although the chemical outcomes in this case will be rather different. METALLIC STRUCTURES The structure of metals The arrangement of the atoms Metals are giant structures of atoms held together by metallic bonds. "Giant" implies that large but variable numbers of atoms are involved - depending on the size of the bit of metal. 12-co-ordination Most metals are close packed - that is, they fit as many atoms as possible into the available volume. Each atom in the structure has 12 touching neighbours. Such a metal is described as 12-co-ordinated. Each atom has 6 other atoms touching it in each layer. There are also 3 atoms touching any particular atom in the layer above and another 3 in the layer underneath.

This second diagram shows the layer immediately above the first layer. There will be a corresponding layer underneath. (There are actually two different ways of placing the third layer in a close packed structure). 8-co-ordination Some metals (notably those in Group 1 of the Periodic Table) are packed less efficiently, having only 8 touching neighbours. These are 8-co-ordinated. The left hand diagram shows that no atoms are touching each other within a particular layer. They are only touched by the atoms in the layers above and below. The right hand diagram shows the 8 atoms (4 above and 4 below) touching the darker coloured one. Dislocations It would be misleading to suppose that all the atoms in a piece of metal are arranged in a regular way. Any piece of metal is made up of a large number of "crystal grains", which are regions of perfect regularity. At the grain boundaries atoms have become misaligned. The grain boundaries are also known as dislocations.

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The physical properties of metals •

Melting points and boiling points Metals tend to have high melting and boiling points because of the strength of the metallic bond. The strength of the bond varies from metal to metal and depends on the number of electrons which each atom delocalises into the sea of electrons, and on the packing. Group 1 metals like sodium and potassium have relatively low melting and boiling points mainly because each atom only has one electron to contribute to the bond - but there are other problems as well: a- Group 1 elements are also inefficiently packed (8-co-ordinated), so that they aren't forming as many bonds as most metals. b- They have relatively large atoms (meaning that the nuclei are some distance from the delocalised electrons) which also weakens the bond. • Electrical conductivity Metals conduct electricity. The delocalised electrons are free to move throughout the structure in 3-dimensions. They can cross grain boundaries. Even though the pattern may be disrupted at the boundary, as long as atoms are touching each other, the metallic bond is still present. Liquid metals also conduct electricity, showing that although the metal atoms may be free to move, the delocalisation remains in force until the metal boils. • Thermal conductivity Metals are good conductors of heat. Heat energy is picked up by the electrons as additional kinetic energy (it makes them move faster). The energy is transferred throughout the rest of the metal by the moving electrons. • Malleability and ductility Metals are described as malleable (can be beaten into sheets) and ductile (can be pulled out into wires). This is because of the ability of the atoms to roll over each other into new positions without breaking the metallic bond. If a small stress is put onto the metal, the layers of atoms will start to roll over each other. If the stress is released again, they will fall back to their original positions. Under these circumstances, the metal is said to be elastic.

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If a larger stress is put on, the atoms roll over each other into a new position, and the metal is permanently changed. Hardness: This rolling of layers of atoms over each other is hindered by grain boundaries because the rows of atoms don't line up properly. It follows that the more grain boundaries there are (the smaller the individual crystal grains), the harder the metal becomes. Offsetting this, because the grain boundaries are areas where the atoms aren't in such good contact with each other, metals tend to fracture at grain boundaries. Increasing the number of grain boundaries not only makes the metal harder, but also makes it more brittle. Controlling the size of the crystal grains If you have a pure piece of metal, you can control the size of the grains by heat treatment or by working the metal. Heating a metal tends to shake the atoms into a more regular arrangement 4

- decreasing the number of grain boundaries, and so making the metal softer. Banging the metal around when it is cold tends to produce lots of small grains. Cold working therefore makes a metal harder. To restore its workability, you would need to reheat it. You can also break up the regular arrangement of the atoms by inserting atoms of a slightly different size into the structure. Alloys such as brass (a mixture of copper and zinc) are harder than the original metals because the irregularity in the structure helps to stop rows of atoms from slipping over each other. GIANT COVALENT STRUCTURES The physical properties of diamond Diamond: •

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has a very high melting point (almost 4000°C). Very strong carbon-carbon covalent bonds have to be broken throughout the structure before melting occurs. is very hard. This is again due to the need to break very strong covalent bonds operating in 3-dimensions. doesn't conduct electricity. All the electrons are held tightly between the atoms, and aren't free to move. is insoluble in water and organic solvents. There are no possible attractions which could occur between solvent molecules and carbon atoms which could outweigh the attractions between the covalently bound carbon atoms.

The atoms within a sheet are held together by strong covalent bonds - stronger, in fact, than in diamond because of the additional bonding caused by the delocalised electrons. So what holds the sheets together? In graphite you have the ultimate example of van der Waals dispersion forces (to be explained later). As the delocalised electrons move around in the sheet, very large temporary dipoles can be set up, which will induce opposite dipoles in the sheets above and below - and so on throughout the whole graphite crystal. The physical properties of graphite Graphite: •

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has a high melting point, similar to that of diamond. In order to melt graphite, it isn't enough to loosen one sheet from another. You have to break the covalent bonding throughout the whole structure. has a soft, slippery feel, and is used in pencils and as a dry lubricant for things like locks. You can think of graphite rather like a pack of cards - each card is strong, but the cards will slide over each other, or even fall off the pack altogether. When you use a pencil, sheets are rubbed off and stick to the paper. has a lower density than diamond. This is because of the relatively large amount of space that is "wasted" between the sheets. is insoluble in water and organic solvents - for the same reason that diamond is insoluble. Attractions between solvent molecules and carbon atoms will never be strong enough to overcome the strong covalent bonds in graphite.

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conducts electricity. The delocalised electrons are free to move throughout the sheets. If a piece of graphite is connected into a circuit, electrons can fall off one end of the sheet and be replaced with new ones at the other end.

The physical properties of silicon dioxide Silicon dioxide: •

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It has a high melting point - varying depending on what the particular structure is (remember that the structure given is only one of three possible structures), but around 1700°C. Very strong silicon-oxygen covalent bonds have to be broken throughout the structure before melting occurs. It is hard. This is due to the need to break the very strong covalent bonds. It doesn't conduct electricity. There aren't any delocalised electrons. All the electrons are held tightly between the atoms, and aren't free to move. It is insoluble in water and organic solvents. There are no possible attractions which could occur between solvent molecules and the silicon or oxygen atoms which could overcome the covalent bonds in the giant structure.

MOLECULAR STRUCTURES The physical properties of molecular substances Molecules are made of fixed numbers of atoms joined together by covalent bonds, and can range from the very small (even down to single atoms, as in the noble gases) to the very large (as in polymers, proteins or even DNA). The covalent bonds holding the molecules together are very strong, but these are largely irrelevant to the physical properties of the substance. Physical properties are governed by the intermolecular forces - forces attracting one molecule to its neighbours - van der Waals attractions or hydrogen bonds. •

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Melting and boiling points Molecular substances tend to be gases, liquids or low melting point solids, because the intermolecular forces of attraction are comparatively weak. You don't have to break any covalent bonds in order to melt or boil a molecular substance. The melting or boiling point will depend on the strength of the intermolecular forces. Solubility in water Most molecular substances are insoluble (or only very sparingly soluble) in water. Those which do dissolve often react with the water, or else are capable of forming hydrogen bonds with the water. Solubility in organic solvents Molecular substances are often soluble in organic solvents - which are themselves molecular. Both the solute (the substance which is dissolving) and the solvent are likely to have molecules attracted to each other by weak intermolecular forces. Although these attractions will be disrupted when they mix, they are replaced by similar ones between the two different sorts of molecules. Electrical conductivity Molecular substances won't conduct electricity. Even in cases where electrons may be delocalised within a particular molecule, there isn't sufficient contact between the molecules to allow the electrons to jump and move through the whole solid or liquid.

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INTERMOLECULAR BONDING - VAN DER WAALS FORCES So what are intermolecular attractions? Intermolecular attractions are attractions between one molecule and a neighbouring molecule. All molecules experience intermolecular attractions, although in some cases those attractions are very weak. Even in a gas like hydrogen, H 2, if you slow the molecules down by cooling the gas, the attractions are large enough for the molecules to stick together eventually to form a liquid and then a solid. These forces are generally known as van der Waals forces In hydrogen's case the attractions are so weak that the molecules have to be cooled to 21 K (-252°C) before the attractions are enough to condense the hydrogen as a liquid. Helium's intermolecular attractions are even weaker - the molecules won't stick together to form a liquid until the temperature drops to 4 K (-269°C). Dispersion Forces: a kind of van der Waals interactions Dispersion forces (one of the two types of van der Waals force we are dealing with) are also known as "London forces". The origin of dispersion forces: Temporary fluctuating dipoles Attractions are electrical in nature. In a symmetrical molecule like hydrogen, however, there doesn't seem to be any electrical distortion to produce positive or negative parts. But that's only true on average. The elliptical diagram represents a small symmetrical molecule - H2, perhaps, or Br2. The even shading shows that on average there is no electrical distortion. But the electrons are mobile, and at any one instant they might find themselves towards one end of the molecule, making that end δ-. The other end will be temporarily short of electrons and so becomes δ+. An instant later the electrons may well have moved up to the other end, reversing the polarity of the molecule. (See figure) This constant "sloshing around" of the electrons in the molecule causes rapidly fluctuating dipoles even in the most symmetrical molecule. It even happens in monatomic molecules - molecules of noble gases, like helium, which consist of a single atom. If both the helium electrons happen to be on one side of the atom at the same time, the nucleus is no longer properly covered by electrons for that instant. How temporary dipoles give rise to intermolecular attractions Imagine a molecule which has a temporary polarity being approached by one which happens to be entirely non-polar just at that moment. (A pretty unlikely event) As the right hand molecule approaches, its electrons will tend to be attracted by the slightly positive end of the left hand one. This sets up an induced dipole in the approaching molecule, which is orientated in such a way that the δ+ end of one is

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attracted to the δ- end of the other. An instant later the electrons in the left hand molecule may well have moved up the other end. In doing so, they will repel the electrons in the right hand one.

The polarity of both molecules reverses, but you still have δ+ attracting δ-. As long as the molecules stay close to each other the polarities will continue to fluctuate in synchronisation so that the weak attraction is always maintained. There is no reason why this has to be restricted to two molecules. As long as the molecules are close together this synchronised movement of the electrons can occur over huge numbers of molecules. This diagram shows how a whole lattice of molecules could be held together in a solid using van der Waals dispersion forces. An instant later, of course, you would have to draw a quite different arrangement of the distribution of the electrons as they shifted around but always in synchronisation. The strength of dispersion forces between molecules is much weaker than the covalent bonds within molecules. It isn't possible to give any exact value, because the size of the attraction varies considerably with the size of the molecule and its shape. Molecular size affects the strength of the dispersion forces The boiling points of the noble gases are

helium neon

269°C 246°C 186°C

krypton -152°C xenon

-108°C

radon -62°C All of these elements have monatomic argon molecules. The reason that the boiling points increase as you go down the group is that the number of electrons increases, and so also does the radius of the atom. The more electrons you have and the more distance over which they can move, the bigger the possible temporary dipoles and therefore the bigger the dispersion forces. This is the reason that (all other things being equal) bigger molecules have higher boiling points than small ones. Bigger molecules have more electrons and more distance over which temporary dipoles can develop - and so the bigger molecules are "stickier".

Molecular shape affects the strength of the dispersion forces The shapes of the molecules also matter. Long thin molecules can develop bigger temporary dipoles due to electron movement than short fat ones containing the same numbers of electrons. Long thin molecules can also lie closer together - these attractions are at their most effective if the molecules are really close.

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For example, the hydrocarbon molecules butane and 2-methylpropane both have a molecular formula C4H10, but the atoms are arranged differently. In butane the carbon atoms are arranged in a single chain, but 2-methylpropane is a shorter chain with a branch. Butane has a higher boiling point because the dispersion forces are greater. The molecules are longer (and so set up bigger temporary dipoles) and can lie closer together than the shorter, fatter 2-methylpropane molecules. Dipole-dipole interactions: a second kind of van der Waals interactions A molecule like HCl has a permanent dipole because chlorine is more electronegative than hydrogen. These permanent, in-built dipoles will cause the molecules to attract each other rather more than they otherwise would if they had to rely only on dispersion forces. It's important to realise that all molecules experience dispersion forces. Dipoledipole interactions are not an alternative to dispersion forces - they occur in addition to them. Molecules which have permanent dipoles will therefore have boiling points rather higher than molecules which only have temporary fluctuating dipoles. Surprisingly dipole-dipole attractions are fairly minor compared with dispersion forces, and their effect can only really be seen if you compare two molecules with the same number of electrons and the same size. For example, the boiling points of ethane, CH3CH3, and fluoromethane, CH3F, are shown in the figure The higher boiling point of fluoromethane is due to the large permanent dipole on the molecule because of the high electronegativity of fluorine. However, even given the large permanent polarity of the molecule, the boiling point has only been increased by some 10°. Here is another example showing the dominance of the dispersion forces. Trichloromethane, CHCl3, is a highly polar molecule because of the electronegativity of the three chlorines. There will be quite strong dipole-dipole attractions between one molecule and its neighbours. On the other hand, tetrachloromethane, CCl4, is nonpolar. The outside of the molecule is uniformly δ- in all directions. CCl 4 has to rely only on dispersion forces. So which has the highest boiling point? CCl4 does, because it is a bigger molecule with more electrons. The increase in the dispersion forces more than compensates for the loss of dipole-dipole interactions. The boiling points are: CHCl3 61.2°C CCl4 76.8°C

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INTERMOLECULAR BONDING - HYDROGEN BONDS The evidence for hydrogen bonding Many elements form compounds with hydrogen - referred to as "hydrides". If you plot the boiling points of the hydrides of the Group 4 elements, you find that the boiling points increase as you go down the group. The increase in boiling point happens because the molecules are getting larger with more electrons, and so van der Waals dispersion forces become greater. If you repeat this exercise with the hydrides of elements in Groups 5, 6 and 7, something odd happens. Although for the most part the trend is exactly the same as in group 4 (for exactly the same reasons), the boiling point of the hydride of the first element in each group is abnormally high. In the cases of NH3, H2O and HF there must be some additional intermolecular forces of attraction, requiring significantly more heat energy to break. These relatively powerful intermolecular forces are described as hydrogen bonds. The origin of hydrogen bonding The molecules which have this extra bonding are: Notice that in each of these molecules: •

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The hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen to acquire a significant amount of positive charge. Each of the elements to which the hydrogen is attached is not only significantly negative, but also has at least one "active" lone pair.

Lone pairs at the 2-level have the electrons contained in a relatively small volume of space which therefore has a high density of negative charge. Lone pairs at higher levels are more diffuse and not so attractive to positive things. Consider two water molecules coming close together. The δ+ hydrogen is so strongly attracted to the lone pair that it is almost as if you were beginning to form a co-ordinate (dative covalent) bond. It doesn't go that far, but the attraction is significantly stronger than an ordinary dipole-dipole interaction. Hydrogen bonds have about a tenth of the strength of an average covalent bond, and are being constantly broken and reformed in liquid water. If you liken the covalent bond

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between the oxygen and hydrogen to a stable marriage, the hydrogen bond has "just good friends" status. On the same scale, van der Waals attractions represent mere passing acquaintances! Hydrogen bonding in alcohols An alcohol is an organic molecule containing an -O-H group. Any molecule which has a hydrogen atom attached directly to an oxygen or a nitrogen is capable of hydrogen bonding. Such molecules will always have higher boiling points than similarly sized molecules which don't have an -O-H or an -N-H group. The hydrogen bonding makes the molecules "stickier", and more heat is necessary to separate them. Ethanol, CH3CH2-O-H, and methoxymethane, CH3-O-CH3, both have the same molecular formula, C2H6O. They have the same number of electrons, and a similar length to the molecule. The van der Waals attractions (both dispersion forces and dipole-dipole attractions) in each will be much the same. However, ethanol has a hydrogen atom attached directly to an oxygen - and that oxygen still has exactly the same two lone pairs as in a water molecule. Hydrogen bonding can occur between ethanol molecules, although not as effectively as in water as we shall see... The hydrogen bonding is limited by the fact that there is only one hydrogen in each ethanol molecule with sufficient δ+ charge. In methoxymethane, the lone pairs on the oxygen are still there, but the hydrogens aren't sufficiently δ+ for hydrogen bonds to form. Except in some rather unusual cases, the hydrogen atom has to be attached directly to the very electronegative element for hydrogen bonding to occur. The boiling points of ethanol and methoxymethane show the dramatic effect that the hydrogen bonding has on the stickiness of the ethanol molecules: ethanol (with hydrogen bonding) 78.5°C methoxymethane (without hydrogen bonding) -24.8°C

The hydrogen bonding in the ethanol has lifted its boiling point about 100°C. It is important to realise that hydrogen bonding exists in addition to van der Waals attractions. For example, all the following molecules contain the same number of electrons, and the first two are much the same length. The higher boiling point of the butan-1-ol is due to the additional hydrogen bonding. Comparing the two alcohols (containing -OH groups), both boiling points are high because of the additional hydrogen bonding due to the hydrogen attached directly to the oxygen - but they aren't the same. The boiling point of the 2-methylpropan-1-ol isn't as high as the butan-1-ol because the branching in the molecule makes the van der Waals attractions less effective than in the longer butan-1-ol.

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Hydrogen bonding in organic molecules containing nitrogen Hydrogen bonding also occurs in organic molecules containing N-H groups - in the same sort of way that it occurs in ammonia. Examples range from simple molecules like CH3NH2 (methylamine) to large molecules like proteins and DNA. The two strands of the famous alpha-helix in DNA are held together by hydrogen bonds involving N-H groups. Although nitrogen has two hydrogen atoms to form hydrogen bonds it has just one pair of electrons to offer. These opposite effects seem to balance each other. But being nitrogen less electronegative than oxygen, hydrogens linked to nitrogen are not as positive as those bonded to oxygen. Hence hydrogen bridges are not as strong (compare ammonia and water) Water as a "perfect" example of hydrogen bonding Notice that each water molecule can potentially form four hydrogen bonds with surrounding water molecules. There are exactly the right numbers of δ+ hydrogens and lone pairs so that every one of them can be involved in hydrogen bonding. This is why the boiling point of water is higher than that of ammonia or hydrogen fluoride. In the case of ammonia, the amount of hydrogen bonding is limited by the fact that each nitrogen only has one lone pair. In a group of ammonia molecules, there aren't enough lone pairs to go around satisfying all the hydrogens. In hydrogen fluoride, the problem is a shortage of hydrogens. In water, there is exactly the right number of each. Water could be considered as the "perfect" hydrogen bonded system. Ice There are lots of different ways that the water molecules can be arranged in ice. This is one of them, but NOT the common one. The one below is known as "cubic ice", or "ice Ic". It is based on the water molecules arranged in a diamond structure. This is just a small part of a structure which extends over huge numbers of molecules in three dimensions. In the diagram, the lines represent hydrogen bonds. The lone pairs that the hydrogen atoms are attracted to are left out for clarity. Cubic ice is only stable at temperatures below -80°C. The ice you are familiar with has a different, hexagonal structure. It is called "ice Ih". The unusual density behaviour of water The hydrogen bonding forces a rather open structure on the ice - if you made a model of it, you would find a significant amount of wasted space. When ice melts, the structure breaks down and the molecules tend to fill up this wasted space. This means that the water formed takes up less space than the original ice. Ice is a very unusual solid in this respect - most solids show an increase in volume on melting. When water freezes, the opposite happens - there is an expansion as the hydrogen bonded structure establishes. Most liquids contract on freezing.

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Remnants of the rigid hydrogen bonded structure are still present in very cold liquid water, and don't finally disappear until 4°C. The density of water increases from 0°C to 4°C as the molecules free themselves from the open structure and take up less space. After 4°C, the thermal motion of the molecules causes them to move apart and the density falls. That's the normal behaviour with liquids on heating. Water as a solvent for ionic compounds When an ionic substance dissolves in water, water molecules cluster around the separated ions. This process is called hydration. Water frequently attaches to positive ions by co-ordinate (dative covalent) bonds in which the oxygen atoms act as donors and the metal ions as acceptors. In the case of negative ions water bonds to them through hydrogen bonds. The diagram shows the potential hydrogen bonds formed to a chloride ion, Cl-. Although the lone pairs in the chloride ion are at the 3-level and wouldn't normally be active enough to form hydrogen bonds, in this case they are made more attractive by the full negative charge on the chlorine. However complicated the negative ion, there will always be lone pairs that the hydrogen atoms from the water molecules can hydrogen bond to. Water as a solvent for some covalent substances As we have previously stated most molecular substances are insoluble (or only very sparingly soluble) in water. Those which do dissolve are capable of forming hydrogen bonds with the water. Two different situations: methane and ammonia Methane is water insoluble. Methane (CH4) itself isn't the problem. Methane is a gas, and so its molecules are already separate - water doesn't need to pull them apart from one another. The problem is the strong interaction (hydrogen bonds) between water molecules. If methane were to dissolve, it would have to force its way between water molecules and so break hydrogen bonds. That costs a reasonable amount of energy. The only attractions possible between methane and water molecules are the much weaker van der Waals forces - and not much energy is released when these are set up. It simply isn't energetically profitable for the methane and water to mix. Ammonia (NH3) shows the opposite behaviour: it dissolves in water very easily in a very exothermic (heat releasing) process. Ammonia has the ability to form hydrogen bonds. When the hydrogen bonds between water molecules are broken, they can be replaced by equivalent bonds between water and ammonia molecules. Other common substances which are freely soluble in water because they can hydrogen bond with water molecules include ethanol (alcohol) and sucrose (sugar).

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PROBLEMS 1- Using the electronegativity table classify the following substances as having more or less than 50% “ionic character”: NaCl Na2O CO2 NH3 BaCl2 Al F3 Ca3N2. 2- Write the following bonds in order of increasing polarity: C – O H O- H N – H S – O

C – Cl C –

3- Why are the noble gases not shown in the electronegativity table? 4- Explain why KF has a formula but not a molecule. 5- Which of each pair do you expect to have a higher melting point? Explain your choice. (a) MgO or CaS (b) MgO or NaF (c) LiF or BaSe (d) NaF or NaI 6- Explain the following a- Graphite conducts electricity but diamond does not. b- You can write with a graphite pencil. c- Brass is harder than copper. d- Heating treatments makes metals softer. 7- Which of each set do you expect to have a higher melting point? Explain your choice. (a) Li or K (b) Na or Ca (c) K or Fe 8- The following table shows the melting point (in K) of the elements of the 3rd period. Explain the trend.

9- The following table shows the properties of some chemicals found in a lab. According to the information shown: a- ¿Which of the three is propanone (C3H6O)? b- ¿Which of them is copper? c- ¿Which is lithium fluoride? Substance Meeting point A B C

1080 ºC 1290 ºC -127 ºC

Electric conductivity Solid Liquid Very good Very good Poor Very good Poor Poor

Water solubility Insoluble Soluble Soluble

Solubility in solvents Insoluble Insoluble Soluble

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10- Say whether the following statement is true or false. Give reasons for your answer. “Argon atoms have a mass of 40 units (20 protons and 20 electrons) and Krypton a mass of 84 units (38 protons and 46 neutrons). Being Kr particles heavier than Ar particles the gravitational force among them is stronger and consequently its boiling point is higher”. (b.p. (Ar) = -185ºC and b.p. (Kr) = -152ºC) 11- Decide which of the members of each pair will show a higher boiling point. Explain your options a- Na Cl b- F2

HCl Cl2

b- KCl

AlCl3

c- CH3-CH2-CH2-OH

CH3-CH2-OH

d- CH3OH

CH3NH2

e- CH3CH2NH2

CH3NHCH3

f- CH3-CH2-O-CH2-CH3

CH3-CH2-CH2-CH2-OH

g-

h-

i-

12- Which factors affect the solubility of molecular substances in water? 13- Which of each pair can be predicted to have a higher solubility in water? a- CH3-CH2-CH2 -CH2-OH

CH3-CH2-OH

b- CH3-CH2-O-CH2-CH3

CH3-CH2-CH2-CH2-OH

c- CH3-CH2-CH2-CH2-CH3

CH3-CH2-O-CH2-CH3

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