ENERGY. Bonding and Structure

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1 Bonding and Structure Bonding is a fundamental concept in Chemistry without it, the subject would not exist. Both the physical and chemical properties of an element or compound depend on the bonding involved, but Why do atoms bond together? The answer to this question will be the answer to most why questions in this subject. ENERGY We know that bond breaking is endothermic, and bond making is exothermic. When two (or many) atoms rearrange their electrons in their search for stability (lower energy), a bond is formed between them.

to explain: Melting temperatures (variation between metals)

2 Metallic Bonding Key ideas: Lattice of cations Sea of delocalised electrons Strong electrostatic forces of attraction You must be able to explain: Melting temperatures (variation between metals) Electrical conductivity Thermal conductivity Malleability and ductility Questions p37

Metallic bonding: Sodium, magnesium and aluminium This type of bonding is strong which explains the relatively high melting and boiling points. What trend do you notice across the metals?

3 Metallic bonding: Sodium, magnesium and aluminium This type of bonding is strong which explains the relatively high melting and boiling points. What trend do you notice across the metals? There are 4 important factors: 1. Atomic radii get smaller, more close packing. 2. Increased nuclear charge 3. Number of valence electrons involved in bonding increases, Na 1, Mg 2 and Al 3 4. Crystal structure the Na packing is cubic, less efficient than that of Mg and Al.

The effect of these factors from Na to Al is a stronger

and a closer and more negative sea of electrons.

ΔHvap all increase. This also explains the big increase in m.

4 The effect of these factors from Na to Al is a stronger attraction between the increasingly positively charged nuclei and a closer and more negative sea of electrons. Hence, the melting point, the boiling point, ΔHa, ΔHfus and ΔHvap all increase. This also explains the big increase in m.pts between Na and Mg and the small increase between Mg and Al.

5 FCC Face centered cubic (FCC) Body centered cubic (BCC) Hexagonal close packed (HCP)

6 Up to now, we have thought of ionic bonds being the transfer of electrons from a metal to a non-metal and the substance formed being held together by electrostatic forces. A covalent bond was the sharing of electrons by non-metals. We will see, however, that what we considered to be 2 different types of bond are in fact the extremes of the same continuum. IONIC BONDING This type of bonding is the result of electrostatic forces between charged particles ions. Ions can be cations or anions. What evidence do we have to prove their existence? Ions can be simple, compound, complex, eg. Mg 2+ NO 3 - [Zn(CN)6] 4-

7 A species with more protons than electrons will have a positive charge, eg A species with fewer protons than electrons will have a negative charge, eg But why would something want to gain or lose electrons? Yes, again, due to questions of energy and stability. Remember the octet rule, noble gas configuration, etc? Well, it still holds, more or less, but we sometimes have to break it.

8 Demo: Migration of ions Questions p41

of attraction between 2 ions, we should think of ionic bonding

9 Giant Ionic Structures What holds a giant ionic lattice together? Although an ionic bond was traditionally used to mean the force of attraction between 2 ions, we should think of ionic bonding as a system of electrostatic forces of attraction (and repulsion!)

10 Which factors affect the strength of the lattice? Size of the ion - this determines how close together the ions can get and therefore the amount of energy required to completely separate them. Size of the charge on the ion - in general, the larger the charge the more energy required to separate the ions NB1. The way that the ions are arranged - the crystal structure - will have an effect NB2. Some ionic substances exhibit covalent character - we will come back to this

11 Trends in Ionic Radii Going down a group. Going across a period. NB1 Isoelectronic species NB2 Ionic radius and coordination number

12 Physical Properties Read pp40-41 Explain why Ionic substances have high melting points Ionic solids are brittle They conduct electricity when molten or in solution They are soluble in polar solvents

13 Remember these? Dot and cross diagrams Abbreviated

14 Covalent Bonding I know that a covalent bond is the strong electrostatic attraction between two nuclei and the shared pair of electrons between them

15 Covalent Bonding sharing electrons in search of stability The octet rule still applies in many cases, but with larger atoms more electrons can be shared by using the 3d orbitals (3rd period onwards). Homework 1.3.1: Give examples of common covalent molecules that contain single, double and triple covalent bonds. Draw dot and cross diagrams and name them.

16 Compounds or polyatomic ions containing sulphur are good examples of breaking the octet rule, SF6 SO3 SO2 SO3 2- SO4 2- Draw dot and cross diagrams and name these species. Try to draw: O3 P4

CO NH4 + H3O + Draw dot and cross diagrams and name these species. How do we represent a dative bond? How about nitric acid?

17 Dative covalent bonds (sometimes called coordinate covalent bonds) are formed when two atoms share two electrons forming a covalent bond between them, but both electrons originate from only one of the atoms, Eg. CO NH4 + H3O + Draw dot and cross diagrams and name these species. How do we represent a dative bond? How about nitric acid? HNO3 Task 1: Find out what is really happening in the nitrate ion with regards to bonding. Task 2: Draw a do and cross diagram for NO. Why is it called a free radical?

18 An excellent example of the formation of dative bonds can be seen in transition metal complexes, Eg. [Ni(H2O)6] 2+ [Fe(CN)6] 4- Note: the species which donates its electrons to the central metal ion is called a ligand. the hydration of metal ions in this way was only thought to happen with transition metals, but other metal ions have been found to form complexes with water, eg. [Mg(H 2 O) 6 ] 2+ [Al(H 2 O) 6 ] 3+ Try Al2O6 Questions p49,50

BUT Look at the following table: What holds the molecule together is the electrostatic force of attraction between the bond

19 The strength of covalent bonds In general, shorter bonds tend to be stronger i.e. They have a greater bond enthalpy. More energy is required to break them and more energy is released when they are formed. BUT Look at the following table: What holds the molecule together is the electrostatic force of attraction between the bond pair(s) of electrons and the nuclei of the atoms, but these nuclei will also repel each other. It is clear above that this repulsion can weaken a bond, in this case... Questions p44

20 Covalent bonds and overlapping orbitals Remember this? Now we take dot and cross diagrams to a new level. Covalent bonds are formed when orbitals in atoms overlap to increase the electron density between the atoms. In simple covalent bonding, the two orbitals that overlap each contain 1 unpaired electron. In dative bonding, a lone pair of electrons in an orbital of one of the species is donated into an empty orbital on the other species to form the covalent bond.

21 The shape of molecules and ions Valence Shell Electron Pair Repulsion (VSEPR) theory. Important points: Generally speaking, valence electrons in molecules are found as either lone pairs (lp) or bond pairs (bp). These pairs of electron will repel each other and therefore separate as much as possible. Lone pairs are closer to the nucleus and have greater charge density. They therefore repel to a greater extent than bond pairs. This affects the angles in the molecules but not the general shape. Note: You must learn these angles!

22 You need to learn the following shapes, the reasoning behind them and at least one example of each. # of lone pair electrons on 'central' atom # of bonding groups (pair electrons) on 'central' atom Examples Molecular Geometry Bond Angle 0 2 linear trigonal planar Bent less than tetrahedral trigonal pyramidal less than bent less than trigonal bipyramidal 90, 120 and seesaw 90, 120 and T-shaped 90 and linear octrahedral 90 and square pyramidal 90 and 180 SS 2 4 square planar 90 and 180

23 Ions and molecules with double and triple bonds Although multiple bonds may have a slightly higher electron density, to determine the shape of a molecule we treat them as if they were single covalent bonds (also with dative bonds). Challenge: Try to draw dot and cross diagrams for each of the following and then work out and represent their shape. CO2 SO2 SO3

24 SO3 2- SO4 2- NO3 - NO2 + CO3 2- [Cu(H2O)6] 2+ NH4 + H3O + NB. Be prepared to be asked to determine, by analogy, the shapes of molecules that you have never come across. Questions p52

25 * Diamond and graphite are allotropes of carbon. Giant molecular structures These are vast 3-D structures of atoms that are linked by covalent bonds. The 3 typical examples are: Each C atom connected to 4 others by sigma bonds Each C atom connected to 3 others by sigma bonds *

Properties of giant molecular structures Poor conductors of electricity because the electrons are not usually free to move as they can in metallic structures (diamond is a relatively good thermal

Due to the strong bonding in all directions in the structure, they are often very hard, strong and will not dissolve in solvents like water (diamonds is used as the leading edge in cutting tools).

26 Properties of giant molecular structures Poor conductors of electricity because the electrons are not usually free to move as they can in metallic structures (diamond is a relatively good thermal conductor, as it transmits vibrations through its structure).* Thermally very stable: high melting and boiling points (m.pt >3800ºC for diamond and graphite). Due to the strong bonding in all directions in the structure, they are often very hard, strong and will not dissolve in solvents like water (diamonds is used as the leading edge in cutting tools). Diamond has a relatively high density (C atom is relatively light) however, compared with crystals of elements with similar atomic weights, diamond is much more dense! This is due to the very close packing of the atoms. Graphite is less dense than diamond. * The high m.pt, good electrical conductivity and inertness of graphite make it perfect for use in electrolysis as an electrode.

27 Graphite is unique, and deserves some extra attention. The electron in the p orbital which is not involved in the σ bonds, forms a π bond with the C atoms around it a delocalised π system is created. The electrons are thus free to move in this system, along the layers (but not between them!). Graphite therefore conducts electricity but not as well as metals. The layers are held together by weak V de W forces, allowing them to slide over each other. Graphite is used as a dry lubricant.

The physical properties of graphite * 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.

28 The physical properties of graphite * 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. * 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 diamond * 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.

29 The physical properties of diamond * 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 lengths of the carbon-carbon bonds in diamond, graphite, and buckyballs have been measured by a technique called x-ray diffraction. In diamond the bond lengths are 1.54 Α (1.54 x m, the A unit represents 1 x 10-10m)

The other allotrope of carbon - Buckminster Fullerene C60 Hexagonal rings like graphite and alternating pentagonal rings to allow curvature of the surface. Others are oval shaped like a rugby ball.

30 The other allotrope of carbon - Buckminster Fullerene C60 Hexagonal rings like graphite and alternating pentagonal rings to allow curvature of the surface. Others are oval shaped like a rugby ball. They are all called Fullerenes or affectionately Bucky Balls They do dissolve in solvents, and although they are solid, their melting points are not that high. They can be converted into continuous tubes to form very strong fibres of 'pipe like' molecules called 'nanotubes'. These are simple molecules, NOT giant covalent structures.

31 Investigate NANOTECHNOLOGY and Graphene 5min presentation in pairs!!!!! TO INCLUDE: Photos of different types of nanotube Properties Current and possible applications

Silicon(IV)Oxide I think that this undervalued and under loved giant

melting point (1,610 C) and boiling point (2,230 C) - insoluble in water -

.very strong covalent bonds that hold the silicon and oxygen atoms in the

32 Silicon(IV)Oxide I think that this undervalued and under loved giant macromolecule deserves a little of out attention - very hard - very high melting point (1,610 C) and boiling point (2,230 C) - insoluble in water - does not conduct electricity. Due to..very strong covalent bonds that hold the silicon and oxygen atoms in the giant covalent structure. Lots of uses, but mainly used to make Silicon Questions p67

33 H Electronegativity Li 0.98 Be 1.57 This is the capacity that an atom has to pull the electrons in a covalent bond* towards it. B 2.04 C 2.55 N 3.04 O 3.44 F 3.98 Na 0.93 Mg 1.31 The most common scale for measuring electronegativity is the Pauling Scale. EN Increases across a period Al 1.61 Si 1.90 P 2.19 S 2.58 Cl 3.16 EN + K 0.82 Ca 1.00 Sc 1.36 Ti 1.54 V 1.63 Cr 1.66 Mn 1.55 Fe 1.83 Co 1.88 Ni 1.91 Cu 2.00 Zn 1.65 Ga 1.81 Ge 2.01 As 2.18 Se 2.55 Br 2.96 Rb 0.82 Sr 0.95 Y 1.22 Zr 1.33 Nb 1.60 Mo 2.16 Te 1.90 Ru 2.20 Rh 2.28 Pd 2.20 Ag 1.93 Cd 1.69 In 1.78 Sn 1.96 Sb 2.05 Te 2.10 I 2.66 Cs 0.79 Ba 0.89 La 1.10 Hf 1.30 Ta 1.50 W 2.36 Re 1.90 * Not to be confused with electron affinity. Os 2.20 Ir 2.20 Pt 2.28 Au 2.54 Hg 2.00 Tl 2.04 Pb 2.33 Bi 2.02 Po 2.00 At 2.20

34 Try to explain the displayed trends: There are two key factors*: 1. The number of protons in the nucleus. 2. The distance of the bond pair of electrons from the nuclei of the atoms involved. - Shorter bond lengths Effects of electronegativity: - Polar bonds are formed. The electrons in the bond are displaced towards the more electronegative atom (B), leading to a slightly positive and a slightly negative charge on each end of the bond. We call this polar bond a dipole. * Screening can also influence EN but to a lesser extent.

35 If B is a lot more electronegative than A the electron pair is dragged right over to B's end of the bond and in effect, A has lost its electron, and B dominates the bond pair. Ions have been formed. This brings us back to the idea that ionic and covalent bonds lie on extremes of the same spectrum. In a pure covalent bond, the electrons are held on average exactly half way between the atoms. Even the sodium in sodium chloride hasn't completely lost control of its electron but it acts as if it were purely ionic. Lithium iodide is "ionic with some covalent character, where the pair of electrons hasn't moved entirely over to the iodine end of the bond. This affects the properties of Lithium iodide, for example, it dissolves in organic solvents like ethanol. NB. Polarity in bonds is often referred to as intermediate bonding as the effect lies between ionic and covalent bonding which are at opposite ends of the continuum. Questions p47

36 A molecule can be described as polar, but the fact that a molecule contains dipoles (polar bonds) does not make it so. The classic example is: Carbon Tetrachloride Although it contains four polar bonds, it is not polar overall. The positive and negative charges are said to cancel out there is no positive and negative end to the molecule. Another, more general, way of conceiving this is by using the ideas of centre of charge. The centres of the positive and negative charges in the above molecule coincide, and the molecule is not polar overall.

The centres of charge are separated leading to the formation of a dipole in the molecule overall (the size of this is measured as a dipole

37 A close relative of CCl4 is a polar molecule: Chloromethane Here, the hydrogen is less EN than the carbon, and a positive and negative end to the molecule are created. The centres of charge are separated leading to the formation of a dipole in the molecule overall (the size of this is measured as a dipole moment, which is the size of charges multiplied by the separation of the centres of charge). Examples: Note: Polar molecules will dissolve in polar solvents (water), so many organic compounds are soluble both in organic solvents (CCl4) and in polar solvents.

38 Write the charges where necessary and decide whether the molecule is polar overall. Note: Lone pairs can affect the strength of polarity of a molecule. Nitrogen trifluoride is less polar than expected, given that ammonia is a very polar molecule. Why? Questions p54

39 Intermolecular forces Covalent bonds are intramolecular forces they are relatively strong and difficult to break in reactions. There are three types of intermolecular forces, that are all related to some kind of dipole: Hydrogen bonding Temporary-induced dipole-dipole interactions: London forces* Dipole-dipole interactions* NB(1) * These two types of forces are known as Van der Waals forces, although some sillabi are confusing. NB(2) London forces are sometimes called dispersion forces.

40 1. Hydrogen bonding Look at the data for the melting points of the following hydrides: Why are the hydrides of fluorine, oxygen and nitrogen anomalous?

More heat energy is required to overcome the intermolecular forces in these compounds (luckily for us!) which implies that additional bonding must exist between the molecules.

41 More heat energy is required to overcome the intermolecular forces in these compounds (luckily for us!) which implies that additional bonding must exist between the molecules. Due to the high EN of F, O and N, there is a large dipole produced when these atoms attach to a hydrogen atom, and the negative charges on one molecule are strongly attracted to the positive charges on an adjacent molecule. Notice that the lone pairs add to the size of the dipole in these molecules

molecules and is less dense than liquid water so ice floats

42 As water freezes, there are many more H-bonds. The open structure of ice consists of rings of 6 water molecules and is less dense than liquid water so ice floats (luckily for fish!). We represent the hydrogen bonds between molecules as dotted lines.

43 Density of water Analyse the graph on the right and discuss the implications Investigate: Triple Point

44 Compounds that can H-bond with water are very soluble in it, even if they are organic, Eg.

The bond angle of H-bonds is often stated to be 180º, but this

45 Note: the importance of H-bonds in proteins and DNA/ RNA to hold the helices together. DNA RNA NB. The bond angle of H-bonds is often stated to be 180º, but this can vary. In simple situations (e water-water, or water-ethanol, it can be shown as follows...

other end temporarily short of electrons and so becomes +. This causes instantaneous temporary dipoles both in atoms eg.

46 2. Temporary (or induced) dipole-dipole interactions: London forces All molecules experience dispersion forces. Electrons are mobile, and at any one instant they might find themselves towards one end of the molecule, making that end and the other end temporarily short of electrons and so becomes +. This causes instantaneous temporary dipoles both in atoms eg. The noble gases, and in molecules eg. The alkanes. This temporary dipole can cause an induced dipole in an adjacent molecule, and from this molecule to others, etc.

47 The boiling points of the noble gases are: As you go down the group is that the number of electrons increases, as 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 London forces. helium neon argon krypton xenon radon -269 C -246 C -186 C -152 C -108 C -62 C

Generally, bigger molecules have higher boiling points as they have more electrons and more distance over which temporary dipoles can develop - bigger molecules are "stickier".

48 Generally, bigger molecules have higher boiling points as they have more electrons and more distance over which temporary dipoles can develop - bigger molecules are "stickier". Another factor is the shape of the molecule. Long thin molecules can develop bigger temporary dipoles and they can also lie closer together making the attractions more effective. They are said to have more surface area in contact with adjacent molecules.

49 3. Dipole-dipole interactions Give examples of molecules that contain a permanent dipole. It is a general misconception in Chemistry that London forces are the weakest type of intermolecular force, but this is not the case. However, when permanent dipoles interact, the small forces of attraction between the opposite charges of the polar molecules add to the intermolecular forces. Consider two molecules with the same number of electrons and very similar sizes. Do the differences in b.pts make sense?

50 CHCl 3 CCl C 76.8 C CCl4 has the higher b.pt, as it is a bigger molecule with more electrons. The larger dispersion forces are more significant than the dipole-dipole interactions. Does the b.pt data for chloromethane and tetrachloromethane make sense?

51 How the type of bonding affects properties The key (new?) idea to grasp is the importance and the effect of the intra- and intermolecular forces in a substance. IONIC compounds are held together by the strong electrostatic forces between the ions. MOLECULAR COVALENT compounds have strong intramolecular forces (covalent bonds) but their physical properties are determined by the weaker intermolecular forces between the individual molecules (H-bonding, London forces and dipole-dipole interactions). Questions p58

propane Nb. Alcohols are much less volatile (higher bpt.

52 Question: The following substances have a similar number of electrons. Compare and contrast their boiling points: ethanol, sodium chloride, propane Nb. Alcohols are much less volatile (higher bpt. than other organic molecules with similar Mrs due to hydrogen bonding Questions p64

53 Back to this graph You need to be able to explain the trends in the boiling temperatures of the hydrogen halides (or the other halides mentioned, perhaps).

54 Solubility described... Polar substances will generally dissolve in polar solvents, eg. KCl in water. NB. This includes polar organic substances such as glucose, etc. Polar substances will not usually dissolve in non-polar solvents, eg. KCl in pentane Non-polar substances will dissolve in non-polar solvents, but not in polar solvents Polar solvents mix completely with one another, as do non-polar solvents they are miscible Separate layers are formed when we mix polar and non polar solvents they are immiscible. The least dense is the top layer. Organic molecules that contain dipoles will dissolve in solvents with similar characteristics, however, even though many organic molecules are polar due to the dipole that they contain, they will not dissolve in water as they are unable to hydrogen bond with water, e.g.. the halogenaoalkanes

55 Solubility explained Ionic substances in water, Lattice Enthalpies and Hydration Enthalpies In order for an ionic compound to dissolve, the lattice must be broken down (endothermic) and then the resulting ions are hydrated (hydration enthalpy - exothermic). Solubility is favoured if the energy released in the latter is greater that that required in the former, i.e. the overall process is exothermic. This explains the temperature change when we dissolve ionic substances

ethanol molecules. However, solubility falls as the length of the hydrocarbon chain and you may well end up with two layers in your test tube. NB. 1.

56 Solubility of small alcohols in water The small alcohols are completely soluble in water. In both pure water and pure ethanol the main intermolecular attractions are hydrogen bonds, when the molecules are mixed, hydrogen bonds are made between water molecules and ethanol molecules. However, solubility falls as the length of the hydrocarbon chain and you may well end up with two layers in your test tube. NB. 1. Non polar molecules and even polar organic molecules such as the halogenoalkanes are insoluble because they cannot disrupt the h-bond system of the water as they have insufficient polar character. 2. Two non-polar liquids will mix because they both bond with London forces Questions p69

ENERGY. Bonding and Structure

ENERGY. Bonding and Structure Bonding and Structure Bonding is a fundamental concept in Chemistry without it, the subject would not exist. Both the physical and chemical properties of an element or compound depend on the bonding involved,