Message-ID: <073310Z08071994@anon.penet.fi> Newsgroups: alt.drugs From: email@example.com (Dalamar) Date: Fri, 8 Jul 1994 07:25:09 UTC Subject: CHEMISTRY: Bonding and Structure In the following file the numbers immediately following an atoms symbol in a chemical formula should be read as subscript eg C2H6 should be read : CCCCCC H H C H H C H H C HHHHHHH 6 C 222 H H 6 C 2 H H 6 6 CCCCCC 22 H H 6 6 2222 66 The mole is a measure of amount of substance in chemistry and is equivalent to 6.02 x 10(raised to the power of 23) particles. Bonding and Structure _____________________ The vast majority of substances which occur freely in nature, or are synthetically manufactured by man, are not comprised of free atoms, but of atoms held joined together by chemical bonds. How and why do atoms form bonds ? Obviously the formation of a bond must be energetically favourable, leading to a minimum of energy ie the product in which the bonds have been formed must be more stable than the individual atoms, otherwise the bonds would not form. To understand what happens in terms of electronic structure when atoms form bonds consider the group 0 elements. These comprise the inert gases helium, neon, argon, krypton, xenon and radon, all of which are noted for their extreme lack of chemical properties and unreactivity. Atoms of the noble gases do not normally react with any other atoms, so that the gases consist of atoms alone. This lack of reactivity and the fact that the gases are comprised of lone atoms indicates that these atoms are extremely stable, their energy being at such a favourable minimum that it cannot be improved by bond formation. The inert gases all have one thing in common - a complete outer shell of electrons, so we conclude that this is a very stable arrangement. The electrons contained in the outermost shell of an atom are generally the ones concerned with bonding and the formation of _compounds_. When two or more different elements are combined together, so that their atoms become bonded, the resultant substance is called a compound. The properties of the compound usually differs radically from the elements which combined together to form it. A classic example is the formation of water from the elements hydrogen and oxygen. When hydrogen and oxygen are mixed in the correct proportions and a spark or flame applied, a violent reaction occurs in which the hydrogen and oxygen react together to form water. Both oxygen and hydrogen are gases at room temperature, but the product of their reaction together is a clear liquid, without which life would not exist. When atoms form bonds they do so in such a way as to attain a stable electronic configuration. As we have already shown, the most stable configuration is that of a complete outer shell of electrons. There are three ways in which atoms may obtain a stable electronic configuration : by losing, gaining or sharing electrons. If we divide the elements into (a) electropositive elements, whose atoms compete poorly for electrons and give up one or more electrons fairly readily (low ionisation energy), (b) electronegative elements, whose atoms attract electrons strongly and also readily take up electrons, then the following rule of thumb applies : Electropositive element + Electronegative element = Ionic Bond Electronegative element + Electronegative element = Covalent Bond Electropositive element + Electropositive element = Metallic Bond The three modes of bonding described above are : 1. The Ionic Bond. The _ionic bond_ is formed when electrons are transferred from one atom to another, generating cations and anions which are held together by the pure electrostatic attraction of the resulting positive and negative charges. Compounds such as sodium chloride (NaCl), iron sulphide (FeS) and magnesium oxide (MgO) contain this type of bonding. 2. The Covalent bond. The _covalent bond_ is formed by the mutual sharing of electrons between two atoms. Each atom achieves a stable configuration by gaining a share of a number of electrons from the outermost shell of the other atom. Compounds such as methane (CH4), chloroform (CHCl3), hydrogen chloride (HCl) and benzene (C6H6) contain this type of bonding. 3. The metallic bond. This type of bonding, as the name suggests, occurs in metals. The outermost electrons of the metal become _delocalised_, that is they are not associated with any one particular atom, but are free to move from atom to atom in the metal crystal. The structure can then be imagined as an array of metal cations surrounded by a delocalised 'sea' of electrons which hold the cations together. The outstanding electrical conductivity of metals is due to the mobility of these electrons through the lattice. Sodium metal consists of an array of Na+ cations (noble gas config. of neon, K2 L8) held together by the delocalised M1 electrons (sodium originally K2 L8 M1). Ionic and covalent bonding is covered in more detail below. The Ionic Bond ______________ Consider sodium, an electropositive element with low ionisation energy and electronic configuration of K2 L8 M1. When sodium reacts with an electronegative element, for example chlorine, the single electron contained in the M shell is readily lost to give Na+ ion, with the stable electronic configuration of neon, K2 L8. Chlorine, which is of high electronegativity (electron attracting), accepts an electron readily to give the _chloride ion_, Cl-, with the stable electronic configuration of argon, K2 L8 M8. By the transfer of only one electron, from sodium to chlorine, each atom is now 'happier' as it has achieved a more stable electron configuration. The millions of Na+ and Cl- ions which are generated during the reaction form themselves into a regular three dimensional cubic lattice, consisting of alternating Na+ and Cl- ions. Each Na+ ion in the lattice is surrounded by 6 Cl- ions, 4 in the same plane, one in the plane above, and one in the plane below. The diagram below shows a small portion of a single plane of Na+ and Cl- ions as they are arranged in sodium chloride. Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ Cl- Na+ This pattern will repeated not only in the same plane, but also in planes stacked above and below. The planes immediately above and below this one will be arranged so that the chloride ions they contain are above and below the sodium ions in this plane. The _coordination number_ of each ion is _six_. The electrostatic attractive forces between the ions are extremely strong, resulting in a rigid crystal structure and a compound which is a solid. The chemical formula for sodium chloride is written as NaCl, which represents the ratio of sodium ions to chloride ions in the compound. Because the rest of the group I metals (Li, Na, K etc) have similair electronic structure (one electron in outermost shell), they also have similair properties (electropositive, low ionisation energy) and can be expected to react in a similair fasion to sodium with chlorine, or any of the other of the group VII elements (commonly known as the halogens, F, Cl, Br etc), which are all one electron short of an inert gas structure. The resultant compounds will be of the general formula MX, where M represents an alkali metal and X a halogen. Some examples are sodium fluoride (NaF), lithium chloride (LiCl) and potassium iodide (KI). The group II elements are also electropositive and are collectively known as the alkaline earth metals. All of the metals in this group contain 2 electrons in the outermost shell of their atoms, for example the electronic structure of magnesium is K2 L8 M2. In combining with a halogen, an ionic compound of general formula MX2 is formed, where M represents an alkaline earth metal and X a halogen. To obtain an inert gas structure each metal atom must lose 2 electrons. However, each halogen atom requires but one electron to complete its outermost shell, therefore for each M(2+) cation formed there are two X(-) ions also formed, giving a chemical formula of MX2. Examples are magnesium chloride (MgCl2) and calcium fluoride (CaF2). Oxygen is another very electronegative element and with the electronic structure K2 L6, an oxygen atom is two electrons short of attaining the inert gas structure of neon (K2 L8). In compounds with the group I or group II metals, oxygen can accept two electrons to form the _oxide ion_, O(2-), which now has the inert gas structure of neon. Each group I metal atom donates only one electron, therefore the resulting _group I oxides_, have the general formula M2O eg. sodium oxide (Na2O). Each group II metal donates two electrons, giving a general formula of MO for the _group II oxides_, eg. magnesium oxide (MgO). The bonding in these oxides is again ionic (e.pos element + e.neg element). Most of the oxides, although stable, must be prepared by indirect methods as combustion in air gives other products such as peroxides and superoxides. The amount of energy released when one mole of an ionic compound is formed from its constituent ions is known as the _lattice energy_. This figure is usually quite high (eg approx 750 kJ/mol for NaCl) and depends on the nature of the ions and which type of structure they adopt. As well as the NaCl type of lattice which most of the group I halides adopt, many other geometries are formed by other ionic compounds. The reason why any particular geometry is adopted is that the lattice energy is at its most favourable. The Covalent Bond _________________ When two electronegative elements react together, ionic bonds are not formed because both atoms have a tendency to gain electrons. However, both atoms may still achieve an inert gas structure by the mutual sharing of electrons. Consider the element chlorine, which has seven electrons in the outermost shell of its atoms. Chlorine exists under normal conditions as a yellow gas composed of discrete Cl2 molecules. Now consider how two chlorine atoms will combine to form a chlorine molecule (Cl2). If each atom gives a _share_ of one of its outermost electrons to the other, each achieves a full outer shell. As both chlorine atoms are of identical electronegativity, the pair of electrons which now constitute a covalent bond are shared equally between both atoms. Diagramatically this may be represented: x x x x x x x x x x x x x x x Cl + Cl =====> Cl Cl x x x x x x x x x x x x x Chlorine atoms A chlorine molecule Only the outermost electrons are shown in the diagram (the M shell). Each chlorine atom in the chlorine molecule has in its outermost shell six electrons which fully belong to it, plus a share in two more electrons, making a stable octet (inert gas structure of argon, K2 L8 M8) around each atom. A single covalent bond is therefore made up of a shared _pair_ of electrons. A carbon atom is four electrons short of a complete outer shell, therefore it will need to share four electrons and form four bonds. For example, a molecule of carbon tetrachloride is composed of one carbon atom bonded to four chlorine atoms, CCl4. Each chlorine atom is only one electron short of a complete outer shell, so each Cl atom forms only one bond. Diagramatically this may be represented: x x x x Cl x x x x x x x x x x x x x x x x x x C x + 4 Cl ======> Cl C Cl x x x x x x x x x x x x x x x x Cl x x x x Only the outer shell of electrons is shown for each atom. By sharing electrons in this way, both the carbon and all four chlorine atoms attain an inert gas structure. Although these equations and diagrams help us to rationalise the bonding in CCl4, it does not neccessarily follow that the atoms will react directly together. In the case of CCl4, carbon and chlorine do not react directly to CCl4 and carbon tetrachloride must be prepared by indirect reactions. Nitrogen is three electrons short of attaining an inert gas structure and will therefore form three covalent bonds to other atoms. Ammonia has the chemical formula NH3 and is produced by the direct reaction of hydrogen and nitrogen at high pressures : 3 H2 + N2 = 2 NH3 Hydrogen atoms are one electron short of attaining the inert gas structure of helium (K2). Each H atom is therefore capable of forming one covalent bond, as in ammonia (NH3). H x x x x N H x x x x H For the N atom, only the outer electrons are shown. Notice in the structure for ammonia that there are two electrons on the nitrogen which do not form bonds. These two electrons are known as a _lone pair_ and play an important role in the properties of ammonia and its derivatives. The bond which a pair of electrons form is more usually represented by a straight line joining the two atoms, and a lone pair by two dots next to the atom to which they belong. Thus ammonia can be more neatly represented by H | :N-H The structural formula of the ammonia molecule with its | 3 single covalent bonds between N and H, plus a single H lone pair situated on nitrogen. Each bond line therefore represents a pair of electrons, which can be considered to be in the outer shell of both the atoms it joins. Each H atom has its required 2 electrons in the K shell, the nitrogen has 3 bond pairs, plus its lone pair, making a total of 3x2+2 = 8 electrons in its outermost shell which is the inert gas structure of neon (K2 L8). This is the _structural formula_ of ammonia and shows us the order in which the atoms are connected. The _molecular formula_ for a compound shows us which atoms are present and their numbers, but there could be many ways of fitting the atoms together so that each still forms its required number of bonds. Therefore, it is important to have a way of systematically naming all compounds in such a way that the structural formula can be worked out simply from the name. Even though such a system of naming has been in force a long time, some old common names are still in use. Some large molecules, which commonly have very long systematic names are generally referred to by an agreed common name. Compounds which share the same molecular formula, but differ in the way their atoms are connected or spatially arranged, are known as _isomers_. For example ethanol and dimethylether are related as _structural isomers_ because although they share the same molecular formula C2H6O, the way in which the atoms are connected differs : H H H H | | | | H-C-O-C-H H-C-C-O-H Ethanol and Dimethylether | | | | structural formulas. H H H H Dimethylether Ethanol C2H6O C2H6O Two other types of isomerism that are important are known as geometrical and optical. As well as single covalent bonds, double and triple covalent bonds also exist. For a double bond, two pairs of electrons are mutually shared between the atoms and for a triple bond three pairs of electrons are shared. An example of a compound containing a double bond is ethene (old name ethylene), which has the molecular formula C2H4 : H H | | A molecule of ethene. C=C | | H H Each carbon atom requires a share in 4 electrons in order to complete its outer shell. Each H atom supplies one electron to pair with one of carbons electrons. As there are two H atoms connected to each C this uses up 2 of carbons 4 valency electrons. The only way both C atoms can obtain a complete outer shell is to now share both of their 2 remaining electrons with each other, so that each carbon atom gets a share in two electrons which originate from the neighbouring carbon atom. Nitrogen molecules are diatomic (contains two atoms, N2) and contain a triple bond between N atoms. Each N atom contains 5 electrons in the outermost shell, hence a share in 3 more is required to complete the octet and achieve an inert gas structure. If each N atom shares 3 of its 5 valency electrons with its neighbouring N atom, each achieves a stable octet. Each N atom thus retains two electrons (a lone pair) which fully belong to it, plus gets a share in six others (3 from itself, 3 from the other), thereby completing the octet around each atom. x x :N x x N: The N2 molecule, : represents a lone x x pair of electrons situated on each N. Double and triple bonds also occur between atoms of different types and are most important for the period two elements carbon, nitrogen and oxygen. For example, the carbon-oxygen double bond is very important in organic chemistry, where C=O is known as the _carbonyl_ group and is present in many important classes of compound eg. ketones, aldehydes, amides and esters. An oxygen atom contains six electrons in its outermost shell and therefore requires a share in two more to achieve an inert gas structure. A carbon atom requires a share in four electrons, therefore it shares two of its electrons with oxygen, which satisfies the requirements of oxygen. This still leaves the C atom two electrons short of the inert gas structure, which it achieves via bonding to other atoms. The nature of the other atoms attached to the carbonyl group will determine the reactivity and class of compound we have. Some examples are given below. Structural formula Class Name __________________ _____ ____ H | H-C-H | C=O Ketone Propanone (acetone) | H-C-H | H CH3 | C=O Aldehyde Ethanal (acetaldehyde) | H CH3 | C=O | O-CH2-CH3 Ester Ethylacetate H | C=O | N-CH3 Amide Dimethylformamide | CH3 Common names shown in brackets. For the first compound in the table i drew the complete structural formula. However it is possible to shorten this slightly by writing : H | -CH3 to represent -C-H | H and H H | | -CH2-CH3 to represent -C-C-H | | H H The oxygen atom originally has 6 electrons in its outermost shell and shares two of these when forming two single covalent bonds (as in dimethylether) or one double bond (as in the above compounds). This leaves two lone pairs of electrons situated on oxygen, but these can usually be omitted when drawing the formulae for compounds. From the way we have discussed bonding so far, you may have expected a double covalent bond to be twice the strength of a single bond (if we consider the bonds to be between the same atoms). However, this is not the case and the double bond, although much stronger than a single bond, falls short of being twice the strength by a fair amount. To account for this we must go on another step in complexity and consider a more accurate model for the electronic structure of the atom. This i hope to do in another file if there is interest, but for the moment these basic ideas will suffice. The Coordinate Bond ___________________ So far you have seen that a single covalent bond consists of a pair of mutually shared electrons. One electron of the shared pair originated from one atom and the other electron from the other atom. However, there is a mode of bonding termed _coordinate_, or sometimes _dative_ in which the bond pair originates from the _same_ atom. To see how this is possible consider again the ammonia molecule, NH3. The nitrogen atom in ammonia has a lone pair of electrons. Even though the nitrogen atom has achieved its stable octet of outer electrons, it is still possible for further bonding to N to take place via the lone pair. For example, NH3 will react with a proton (H+, a hydrogen cation, formed by the removal of the single K electron from a H atom) to give: H | The positive charge now resides H-N->H on the N atom in NH4(+). | H The lone pair from the N atom gives the newly attached H the inert gas config. of helium (K2) whilst at the same time it maintains the octet around N. Once formed, this coordinate bond is identical to that of a normal covalent bond and all N-H bonds in NH4(+) are in fact identical. The positive charge originally carried by H(+) is transferred to the nitrogen atom and the resultant cation, NH4(+), is known as the ammonium ion. The bond pair in molecules such as F2 and Cl2 is situated between identical atoms, which are of course of identical electronegativity. Hence the electron pair may be considered to be exactly in the middle of the two atoms. If however the atoms which are linked by a covalent bond are of different electronegativity then the electron pair of the bond will be drawn closer to the more electronegative atom. This results in a _polarised_ bond in which the more electronegative atom aquires a slight negative charge (because it hogs the electrons) and the other a slight positive charge (beacuse the electrons are being dragged away from it). This slight charge separation is represented by d+ and d- (the greek letter delta). For example, consider a molecule A-B, in which A is more electronegative than B. The bond becomes polarised in the direction of A : d- d+ A-B The resulting partial positive and negative charges attract each other and in fact strengthen the bond slightly. This electrostatic attraction is no different to that found in ionic compounds, so the above bond could be described as being partly ionic in character. In fact, if we kept increasing the electronegativity of atom A and decreasing that of B the compound AB would become increasingly more ionic as more and more negative charge built up on atom A. When the difference in electronegativity between A and B is great enough the compound will be ionic and consist of a lattice of A- and B+ ions. Then there is the region between the extremes, where the bond could be described as mainly covalent, but with some ionic character, or mainly ionic, but with some covalent character. Methyl lithium (CH3Li) is an example of a class of compounds known as the organometallics, and the bond is about 40% ionic in character due to the extreme polarisation of the C-Li bond : H d-| d+ In methyl lithium the C-Li bond is H-C-Li extremely polarised. | H Reagents such as MeLi (Me short for methyl, -CH3) are versatile reagents in the synthesis of organic molecules, where the carbon skeleton of the molecule usually has to be constructed from smaller molecules by a series of reactions. Hydrogen Bonding ________________ Hydrogen bonding occurs in compounds which contain a hydrogen atom bonded to a strongly electronegative element, most commonly oxygen and nitrogen. The X-H bond (X=O,N etc) is polarised (d-)X-H(d+). The resultant d+ and d- charges become attracted to the d- and d+ charges on another molecule of the compound, with the result that a weak attractive force comes into play between the molecules. If we consider water : O.........H H Hydrogen bonding in water. / \ \ / H H.........O ... = Hydrogen bond. H . . . \ . . . O . O.........H / . / \ / H H H.....O \ H Water has two H atoms bonded to one O atom and both of these H's can take place in H bonding. The positively polarised H atoms in one molecule attract the negatively polarised O atoms of other water molecules and a 3-D network of hydrogen bonds is established. Hydrogen bonding is much weaker than either covalent or ionic and H-bonds can be broken fairly readily. To break the H bonds requires the input of energy (usually by heating). The high boiling point of water is due to hydrogen bonding. The hydrogen bonds in water are broken if the sample is heated enough (eg by boiling) and the water molecules, with enough thermal energy that the H-bonds can no longer hold them together, enter the gas phase. Some examples of other types of compound which contain H-bonds are alcohols, carboxylic acids, amines and amides. Van der waals Forces of Attraction __________________________________ This is an extremely weak force of attraction which operates between the molecules in covalently bonded compounds. The size of the attractive force generally increases with the weight of the molecule. A good illustration of this principle is the trend in the boiling points of the alkanes, which increase with increasing molecular mass. The alkanes are a family of organic compounds which contain only carbon and hydrogen. Methane, CH4, is the lightest of the alkanes and as such the V.D.W forces of attraction between its molecules are extremely weak, hence methane is a gas at room temperature. For the next heavier alkanes ethane (CH3CH3), propane (CH3CH2CH3) and butane (CH3CH2CH2CH3) the V.D.W forces do increase, but not enough to allow the alkane to be a liquid at room temperature. However, the next members pentane and hexane are fairly volatile liquids at room temperature. The boiling point continues to increase with increasing molecular weight. When the molecular weight is high enough, the V.D.W forces between the molecules will have increased enough so that the alkane becomes a low melting point solid (as in candle wax). Hence most covalent compounds are either gases, liquids or low melting point solids (there is an exception to this where in some cases infinite 3-D covalent structures are formed, as opposed to discrete molecules, as in diamond and silica, in these cases the boiling points are abnormally high). Shapes of Simple Covalent Molecules - VSEPR Theory __________________________________________________ The shapes of most simple covalent molecules can be predicted by using the valence shell electron pair repulsion theory. This theory states that the shape of a molecule is related to the number of electron pairs (bond pairs or lone pairs) in the outer shell of the central atom. It is assumed that the electron pairs arrange themselves to be as far apart as possible in order to minimise the repulsive forces between them (negative charges repel). If the distribution of these pairs can be predicted then so can the shape and bond angle. Consider the structure of a gaseous molecule of beryllium fluoride BeF2. In this molecule the central Be atom forms two single covalent bonds, one bond to each fluorine atom. There are therefore 2 bonding pairs of electrons in the valence shell of the Be atom in BeF2. These 2 pairs will arrange themselves to be as far apart as possible - and this is 180 degrees to each other. The BeF2 molecule is therefore linear, with a F-Be-F bond angle of 180 degrees. You may have noticed that the central Be atom has only 4 electrons in its outermost shell i.e. it does not have a complete inert gas structure. The molecule is described as being electron deficient. A molecule of boron trifluoride, BF3, has a central B atom covalently bonded to three fluorine atoms by single covalent bonds. The three bond pairs arrange themselves so that repulsion is at a minimum - and this is in a plane triangular shape, with the F-B-F bond angles equal to 120 degrees. The fluorine atoms occupy the corners of an equalateral triangle, with the boron atom in the middle. In methane, CH4, there are four bond pairs of electrons around the central carbon atom. The repulsion is at a minimum if the bond pairs arrange themselves tetrahedrally around the C atom i.e. all H-C-H bond angles are 109 degrees 28 minutes. The hydrogen atoms then occupy the corners of a regular tetrahedron and the CH4 molecule is described as tetrahedral. Ammonia, NH3, has four pairs of electrons around the central N atom. These comprise three bonding pairs (one bond to each H atom) and a lone pair. Because the lone pair is not shared with any other atom it is pulled closer to the N atom than are the bond pairs. This results in the lone pair being more replusive than a bond pair, so the order of repulsion between types is Lone pair - Lone pair > Lone pair - Bond pair > Bond pair - Bond pair In ammonia the 4 pairs are again tetrahedrally distributed, with one of the corners of the tetrahedron occupied by the lone pair. This gives the molecule a pyramidal shape: " | Molecule of ammonia. N /|\ H H H The extra repulsion of the lone pair pushes the bonding pairs closer together and thus reduces the H-N-H bond angle from the expected 109 degrees for a regular tetrahedron, to ##### degrees. It is hard to draw 3D diagrams on this terminal - the three H's are not in the plane of the screen! The N forms the apex of a pyramid. Water has four pairs of electrons around the central oxygen atom. These comprise two bond pairs and two lone pairs. Again the distribution of the pairs is roughly tetrahedral, but this time two of the corners of the tetrahedron are occupied by lone pairs. Because there are two lone pairs which provide extra repulsion, the H-O-H bond angle is reduced to #### degrees. The molecule is V-shaped: O / \ H H Molecules with five bond pairs (and no lone pairs) usually adopt a trigonal bipyramid structure eg PCl5 (in the gas phase): * Cl \| P-* /| * Cl Three of the Cl atoms are in the same plane and form an equalateral triangle. These i have represented by a * instead of a Cl. The Cl-P-Cl bond angle (*-P-*) is 120 degrees. The other two chlorine atoms are arranged 180 degrees to each other and at 90 degrees to the plane of the triangle formed by the three Cl's marked *. Three different Cl-P-Cl bond angles are therefore present. Dalamar. ------------------------------------------------------------------------- To find out more about the anon service, send mail to firstname.lastname@example.org. Due to the double-blind, any mail replies to this message will be anonymized, and an anonymous id will be allocated automatically. You have been warned. Please report any problems, inappropriate use etc. to email@example.com. ============================================================================= Message-ID: <104313Z09071994@anon.penet.fi> Newsgroups: alt.drugs From: firstname.lastname@example.org (Dalamar) Date: Sat, 9 Jul 1994 10:39:46 UTC Subject: CHEMISTRY: Bonding and Structure [missing bond angles] Whoops ! When i was writing the file i left the two bond angles for NH3 and H20 blank because i couldn't remember the exact figures. I meant to go and look them up but it must have slipped my mind. Anyway, here they are : NH3 = 106 degrees, 45 minutes H20 = 104 degrees, 27 minutes Dalamar.