Topic 3 Bonding Textbook PDF

TOPIC 3 BONDING AND STRUCTURE A IONIC BONDING | B COVALENT BONDING | C SHAPES OF MOLECULES | D METALLIC BONDING | E SOLID LATTICES Now that we know something of the structure of atoms, in particular their electronic configurations, we can look at how atoms combine. Knowi...

TOPIC 3 BONDING AND STRUCTURE A IONIC BONDING | B COVALENT BONDING | C SHAPES OF MOLECULES | D METALLIC BONDING | E SOLID LATTICES Now that we know something of the structure of atoms, in particular their electronic configurations, we can look at how atoms combine. Knowing about the way substances are bonded, and the effect that the bonding and structure has on chemical and physical properties, is essential to the development of new materials. These new materials are then used to refine and improve products. Computers and mobile phones are smaller, lighter and faster than ever before. New and better materials for clothes and shoes, particularly for use outdoors in bad weather conditions, are constantly being developed. Polymers and composites have almost completely taken the place of more traditional materials such as wood and metal for many uses. Knowledge of the shapes of molecules is of fundamental importance in understanding how medicines work and for the future design of new medicines. Knowing the shapes of enzymes is important in the development of biochemical catalysts for chemical reactions. The shapes of macromolecules such as DNA and proteins are extremely complicated. However, the rules that govern their shapes are the same as the rules used to predict the shapes of simple molecules such as methane, water and ammonia. MATHS SKILLS FOR THIS TOPIC Use angles and shapes in regular 2D and 3D structures Visualise and represent 2D and 3D forms including two-dimensional representations of 3D objects Understand the symmetry of 2D and 3D shapes M03_IASL_CHEM_44860_TPC3_064-097.indd 64 01/06/2018 12:05  1 VELOCITY AND ACCELERATION 1A MOTION 65 SPECIFICATION REFERENCE 1A 1 VELOCITY AND ACCELERATION 1.3.1 What prior knowledge do I need? What will I study in this topic? Metallic, ionic and covalent bonding The nature of metallic, ionic, covalent, polar covalent and dative covalent bonding Using dot-and-cross diagrams to represent ions and molecules The nature of intermolecular interactions, including hydrogen bonding The physical properties of metals, ionic compounds and covalent compounds, both simple The shapes of discrete (simple) molecules molecular and giant structures Electronegativity and polarity of molecules The electronic configurations of the first 36 An explanation of the physical properties of elements in the Periodic Table substances based on their bonding and structure What will I study later? Topic 4 The existence of isomerism in organic compounds Topics 4 and 5 The mechanisms of some reactions of alkanes, alkenes, halogenoalkanes and alcohols Topic 8 Trends in the properties of Group 2 and Group 7 elements Topics 15 and 19 (Book 2: IAL) The mechanisms of some reactions of carbonyls, carboxylic acids, arenes and organic nitrogen compounds Topic 17 (Book 2: IAL) The nature of the bonding in, and the shapes of, transition metal complexes The characteristic properties of transition metals, e.g. the ability to have more than one oxidation state and the ability to act as catalysts M03_IASL_CHEM_44860_TPC3_064-097.indd 65 01/06/2018 12:05  SPECIFICATION 3A 1 THE NATURE OF IONIC REFERENCE 3.1 3.2 3.3 3.5 PART BONDING A dot-and-cross diagram for the reaction between magnesium and LEARNING OBJECTIVES oxygen is shown in fig B. ◼ Describe the formation of ions in terms of loss or Mg 21 22 Mg O gain of electrons. O O ◼ Draw dot-and-cross diagrams to show electrons in Mg 21 22 Mg O cations and anions. ◼ Know that ionic bonding is the result of strong net ▲ fig B Dot-and-cross diagram showing the formation of magnesium and electrostatic attraction between oppositely charged ions. oxide ions. ◼ Know and be able to interpret evidence for the existence of ions using electron density maps and THE NATURE OF IONIC BONDING from the migrations of ions. Ionic bonding occurs in solid materials consisting of a regular array of oppositely charged ions extending throughout a giant lattice network. THE FORMATION OF CATIONS AND ANIONS The most familiar ionic compound is sodium chloride, NaCl. Some ionic compounds can be formed by the direct combination It consists of a regular array of sodium ions, Na+, and chloride of two elements. ions, Cl−, as shown in fig C. FORMATION OF SODIUM AND CHLORIDE IONS For example, sodium chloride can be formed by burning sodium in chlorine: 5 Na1 2Na(s) + Cl2(g) → 2NaCl(s) We can represent the reaction that occurs by two ionic half-equations: 5 Cl2 2Na → 2Na+ + 2e− and Cl2 + 2e− → 2Cl− Each sodium atom has lost one electron to become a positive sodium ion. The chlorine molecule has gained two electrons to become two chloride ions. We can represent the electronic changes involved by dot-and- cross diagrams. 1 2 Na Cl Na 5 Na1 Cl Cl 1 2 Na Na Cl 5 Cl2 ▲ fig A Dot-and-cross diagram showing the formation of sodium and chloride ions. ▲ fig C Structure of sodium chloride. The diagram on the left is an ‘exploded’ version of the structure, LEARNING TIP which is often drawn for the sake of clarity. In practice, the ions It is important when drawing a dot-and-cross diagram to represent are touching one another, as shown in the diagram on the right. the electrons of one atom using a cross and the other using a dot. You only need to show the outer electrons. EXAM HINT It is worth practising drawing a section of an ionic crystal as you may FORMATION OF MAGNESIUM AND OXIDE IONS be asked to do so in an exam. Remember to include a key to show Here is the equation for the formation of magnesium oxide: which ion is which. 2Mg(s) + O2(g) → 2MgO(s) In an ionic solid, there are strong electrostatic interactions between the ions. The ions are arranged in such a way that the electrostatic M03_IASL_CHEM_44860_TPC3_064-097.indd 66 01/06/2018 12:05  TOPIC 3 3A.1 THE NATURE OF IONIC BONDING 67 attractions between the oppositely charged ions are greater than the electrostatic repulsions between ions LEARNING TIP with the same charge. The electrostatic interaction between ions is not directional: all that matters is the distance between two ions, not their orientation with respect to one another. (Compare this with covalent Avoid saying that there is an bonding in Topic 3B.) ‘ionic bond’ between two ions. This is because in an ionic solid, When ions are present, the electrostatic interaction between them tends to be dominant. However, it each ion interacts with many is possible for there to be significant covalent interactions between ions, so you should think of pure other ions, of both the same ionic bonding as an idealised bonding situation. We will develop this concept further in Topic 12B and opposite charge to itself. (Book 2: IAL). The energy that binds the structure does not come from THE STRENGTH OF IONIC BONDING single interactions between ions You can determine the strength of ionic bonding by calculating the amount of energy required in (so-called ionic bonds), but one mole of solid to separate the ions to infinity (i.e. in the gas phase). When they are at an infinite from the interactions between distance from one another, the ions can no longer interact. all of the ions in the lattice. This energy is called the ‘lattice Table A below shows the energy required to separate to infinity the ions in one mole of various energy’ and you will meet this in alkali metal halides. Topic 12B (Book 2: IAL). AMOUNT OF ENERGY REQUIRED TO SEPARATE THE IONS TO INFINITY / kJ mol−1 F− Cl− Br− I− Li+ 1031 848 803 759 + Na 918 780 742 705 + K 817 711 679 651 + Rb 783 685 656 628 table A Energy required to break up a lattice of an ionic compound. For ions of the same charge, the smaller the ions the more energy is required to overcome the electrostatic interactions between the ions and to separate them. The size of the ions is one factor that affects the strength of ionic bonding, which in turn determines how closely packed the ions are in the lattice. The lattice energy for lithium fluoride, Li+F−, is 1031 kJ mol-1. The equivalent energy for magnesium fluoride, Mg2+(F−)2, is 2957 kJ mol−1. The radius of the Mg2+ ion (0.072 nm) is very similar to the radius of the Li+ ion (0.074 nm). The increased charge of the Mg2+ ion compared to the Li+ ion results in a significant increase in the strength of the ionic bonding. When both cation and anion are doubly charged, the energy required to separate the ions is even larger. For magnesium oxide, Mg2+O2−, the value is 3791 kJ mol−1. There is no simple mathematical relationship to describe the effects that ionic radius and ionic charge have on the strength of ionic bonding. The situation is complicated by the way in which the ions pack together to form the lattice, and by the extent to which there are covalent interactions between the ions. In general, however, the smaller the ions and the larger the charge on the ions, the stronger the ionic bonding. EVIDENCE FOR THE EXISTENCE OF IONS Ionic compounds can conduct electricity and undergo electrolysis when either molten or in aqueous solution. This is the most convincing evidence for the existence of ions. For example, when you pass a direct electric current through molten sodium chloride (fig D), sodium is formed at the negative electrode and chlorine is formed at the positive electrode. negative positive electrode d.c. supply electrode Cl2 Na 1 Cl2 Na1 ▲ fig D Electrolysis of molten sodium chloride. M03_IASL_CHEM_44860_TPC3_064-097.indd 67 01/06/2018 12:05  68 3A.1 THE NATURE OF IONIC BONDING TOPIC 3 The explanation for this phenomenon is that: the positive sodium ions migrate towards the negative electrode where they gain electrons and become sodium atoms the negative chloride ions migrate towards the positive electrode where they lose electrons and –ve +ve become chlorine molecules. At the negative electrode: 2Na+ + 2e− → 2Na At the positive electrode: 2Cl− → Cl2 + 2e− Overall equation: 2NaCl → 2Na + Cl2 We can demonstrate the movement of ions by passing a direct current through copper(II) chromate(VI) solution (fig E). Aqueous copper(II) ions, Cu2+(aq), are blue and aqueous chromate(VI) ions, CrO42−(aq), are yellow. ▲ fig E The effect of passing an The Cu2+(aq) ions migrate towards the negative electrode and the solution around this electrode electric current through aqueous turns blue. The CrO42−(aq) ions migrate towards the positive terminal and the solution around this copper(II) chromate. electrode turns yellow. Further evidence for the existence of ions is supplied by electron density maps. Fig F is an electron density map for sodium chloride. A B B A 0.1 nm ▲ fig F Electron density map of sodium chloride produced from X-ray diffraction patterns. The electron density map clearly shows separate ions. A represents a sodium ion and B represents a chloride ion. CHECKPOINT 1. Explain what is meant by the term ‘ionic bonding’. 2. Calcium reacts with fluorine to form the ionic compound calcium fluoride: Ca(s) + F2(g) → CaF2(s) Use a dot-and-cross diagram to show the electronic changes that occur in this reaction. SKILLS REASONING 3. (a) Suggest why the strength of ionic bonding is greater in sodium fluoride than in potassium fluoride. (b) Suggest why the strength of ionic bonding in calcium oxide is approximately four times larger than that in potassium fluoride. S SUBJECT VOCABULARY ionic bonding the electrostatic attraction between oppositely charged ions M03_IASL_CHEM_44860_TPC3_064-097.indd 68 01/06/2018 12:05  1 VELOCITY AND ACCELERATION 1A MOTION 69 SPECIFICATION 3A 2 IONIC RADII AND REFERENCE 3.6 3.7 3.8 3.9 POLARISATION OF IONS LEARNING OBJECTIVES ◼ Understand the effects of ionic radius and ionic charge on the strength of ionic bonding. ◼ Understand reasons for the trends in ionic radii down a group in the Periodic Table, and for a set of isoelectronic ions, e.g. N3− to Al3+. ◼ Understand the meaning of the term polarisation as applied to ions. TRENDS IN IONIC RADII Ionic radii are difficult to measure accurately, and vary according to the environment of the ion. For example, it is important how many oppositely charged ions are touching it (i.e. the co-ordination number). The nature of the ions is also important. There are several different ways of measuring ionic radii and they all produce slightly different values. If you are going to make reliable comparisons using ionic radii, all the values must come from the same source. Remember that there are quite large uncertainties when using ionic radii. Trying to explain things in detail is made difficult because of those uncertainties. GROUP 1 GROUP 7 ION ELECTRONIC IONIC RADIUS / nm ION ELECTRONIC IONIC RADIUS / nm CONFIGURATION CONFIGURATION Li+ 2 0.076 F− 2.8 0.133 + − Na 2.8 0.102 Cl 2.8.8 0.181 + − K 2.8.8 0.138 Br 2.8.18.8 0.196 + − Rb 2.8.18.8 0.152 I 2.8.18.18.8 0.220 table A Trends in ionic radii in Groups 1 and 7. As you go down each group, the ions have more electron shells; therefore, the ions get larger. PERIOD 2 N3− O2− F− PERIOD 3 Na+ Mg2+ Al3+ Number of protons 7 8 9 Number of protons 11 12 13 Electronic configuration 2.8 2.8 2.8 Electronic configuration 2.8 2.8 2.8 DID YOU KNOW? Ionic radius/nm 0.146 0.140 0.133 Ionic radius/nm 0.102 0.072 0.054 When you compare chemical values, you should make table B Trends in ionic radii across a period. sure all the data you are All six of the ions listed in table B are isoelectronic. In other words, they have the same number of comparing come from the electrons and therefore the same electronic configuration. same source. For example, the values in tables B The ionic radius decreases as the number of protons increases. and C have all been taken As the positive charge of the nucleus increases, the electrons are attracted more strongly and are from the Database of Ionic therefore pulled closer to the nucleus. Radii from Imperial College London. However, Chemistry POLARISATION AND POLARISING POWER OF IONS Data Book (JG Stark and HG In an ionic lattice, the positive ion will attract the electrons of the anion. If the electrons are pulled Wallace) gives a value of towards the cation, the anion is polarised since the even distribution of its electron density has been 0.171 for the ionic radius of distorted. the nitride ion, N3−. In either case, the value is larger than The extent to which an anion is polarised by a cation depends on several factors. The two main that for the oxide ion, O2−, as factors are known as Fajan’s rules and are summarised here. expected. M03_IASL_CHEM_44860_TPC3_064-097.indd 69 01/06/2018 12:05  70 3A.2 IONIC RADII AND POLARISATION OF IONS TOPIC 3 Polarisation will be increased by: high charge and small size of the cation (i.e. high charge density of the cation) Cation Anion high charge and large size of the anion. HIGH CHARGE AND SMALL SIZE OF CATIONS The ability of a cation to attract electrons from the anion towards itself is called its polarising power. A cation with a high charge and a small radius has a large polarising power. An approximate value for the polarising power of a cation can be obtained by calculating its charge density. The charge density of a cation is the charge divided by the surface area of the ion. If the ion is assumed Cation Anion to be a sphere, its surface area is equal to 4πr2, where r is the ionic radius. An approximation to the charge density can be determined by dividing the charge by the square of its ionic radius. charge charge density ~ ______ ​​  2 ​​ r Cation Anion HIGH CHARGE AND SMALL SIZE OF ANIONS The ease with which an anion is polarised depends on its charge and its size. Anions with a large charge and a small size are polarised the most easily. In an ionic lattice, the polarisation of the anions creates some degree of sharing of electrons between the two nuclei. That is, some degree of covalent bonding exists. You will learn more about Region where electrons are existing this concept in Topic 12B (Book 2: IAL). in an area of orbital overlap ▲ fig A A representation of a CHECKPOINT cation attracting the electrons of an anion in an ionic lattice. 1. Explain the trend in the following ionic radii: (a) Ca2+ > Mg2+ > Be2+ (b) P3− > S2− > Cl− 2. The table gives the ionic radii of some ions. FORMULA OF ION IONIC RADIUS / nm + Li 0.076 + Na 0.102 2+ Mg 0.072 3+ Al 0.054 Arrange the ions in order of their polarising power. Show how you arrived at your answer. S SUBJECT VOCABULARY polarising power the ability of a positive ion (cation) to distort the electron density of a neighbouring negative ion (anion) polarisation the distortion of the electron density of a negative ion (anion) M03_IASL_CHEM_44860_TPC3_064-097.indd 70 01/06/2018 12:05  1 VELOCITY AND ACCELERATION 1A MOTION 71 SPECIFICATION 3A 3 PHYSICAL PROPERTIES OF REFERENCE 3.1 3.4 PART IONIC COMPOUNDS LEARNING OBJECTIVES DID YOU KNOW? As nearly always in chemistry, there are exceptions. Some ionic ◼ Be able to explain the physical properties of ionic compounds do conduct electricity when solid. For example, solid compounds in terms of their bonding and structure. lithium nitride (Li3N) will conduct electricity and is used in batteries for this reason. Ionic compounds typically have the following physical properties: high melting temperatures SOLUBILITY brittleness Many ionic compounds are soluble in water. We will explain this poor electrical conductivity when solid but good when molten solubility more fully in Topic 12A (Book 2: IAL), including the often soluble in water. part played by entropy changes. HIGH MELTING TEMPERATURES At the moment, you just need to understand that the energy Ionic solids consist of a giant lattice network of oppositely required to break apart the lattice structure and separate the charged ions (see Topic 2B.2). There are many ions in the lattice ions can, in some instances, be supplied by the hydration of and the combined electrostatic forces of attraction among all of the separated ions produced. Both positive and negative ions are the ions is large. attracted to water molecules because of the polarity that water molecules possess (see Topic 3B for an explanation of polarity). A large amount of energy is required to overcome the forces of attraction sufficiently for the ions to break free from the lattice and slide past one another. 5 water molecule BRITTLENESS 2 If a stress is applied to a crystal of an ionic solid, then the layers 1 of ions may slide over one another. stress 5 Na1 The oxygen ends of the water The hydrogen ends of 5 Cl2 molecules are attracted to the water molecules are positive ions. attracted to negative ions. ▲ fig A Effect of stress on an ionic crystal. ▲ fig B Hydration of ions. Ions of the same charge are now side by side and repel one another. The crystals break apart. CHECKPOINT ELECTRICAL CONDUCTIVITY 1. Explain why sodium chloride: Solid ionic compounds do not, in general, conduct electricity. (a) does not conduct electricity when solid, but does when molten This is because there are no delocalised electrons and the ions are also not free to move under the influence of an applied potential (b) has a high melting temperature difference. (c) is soluble in water. However, molten ionic compounds will conduct since the ions are now mobile and will migrate to the electrodes of opposite sign SUBJECT VOCABULARY when a potential difference is applied. If direct current is used, the hydration the process of water molecules being attracted to ions compound will undergo electrolysis as the ions are discharged at in solution and surrounding the ions; the oxygen ends of the water the electrodes. molecules are attracted to the positive ions (cations); the hydrogen ends of the water molecules are attracted to the negative ions EXAM HINT (anions); hydration of ions is an exothermic process (i.e. heat energy Remember, oxidation always takes place at the anode. is released) Aqueous solutions of ionic compounds also conduct electricity and undergo electrolysis, since the lattice breaks down into separate ions when the compound dissolves. M03_IASL_CHEM_44860_TPC3_064-097.indd 71 01/06/2018 12:05  SPECIFICATION REFERENCE 3B 1 COVALENT BONDING 3.10 LEARNING OBJECTIVES ◼ Know that a covalent bond is formed by the overlap of two atomic orbitals each containing a single electron. ◼ Know that a covalent bond is the strong electrostatic attraction between the nuclei of two atoms and the bonding (shared) pair of electrons. ◼ Understand the relationship between bond length and bond strength for covalent bonds. FORMATION OF COVALENT BONDS A covalent bond forms between two atoms when an atomic orbital containing a single electron from one atom overlaps with an atomic orbital, which also contains a single electron, of another atom. The two electrons in the area of overlap are the bonding electrons. They are sometimes referred to as a ‘shared pair of electrons’. The covalent bond is the electrostatic attraction between the two nuclei of the bonded atoms and the pair of electrons shared between them. The atomic orbitals involved can be any of those found in the atoms, but we shall limit our discussion to those involving only s- and p-orbitals. Fig A shows three ways in which these orbitals may overlap. area of overlap area of overlap end on overlap of two end on overlap of two sideways overlap of two s-orbitals (sigma bond) p-orbitals (sigma bond) p-orbitals (pi bond) DID YOU KNOW? ▲ fig A Formation of sigma bonds by end-on overlap of atomic orbitals and a pi bond by sideways overlap of There is a fourth way that p-orbitals. overlap can occur. An s- and p-orbital can overlap end on. An end-on overlap leads to the formation of sigma (σ) bonds. This leads to the formation of a single covalent bond between the two atoms. A sideways overlap of two p-orbitals leads to the formation of a pi (π) bond. A feature of a π bond is that it cannot form until a σ bond has been formed. For this reason π bonds only exist between atoms ▲ fig B Overlap of an s- and that are joined by double or triple bonds. a p-orbital to form a sigma The different types of orbital overlap are shown in the following examples. bond. This overlap, however, can EXAMPLE 1. HYDROGEN only result from atoms of A hydrogen atom has an electronic configuration of 1s1. two different elements. This almost always leads to the When two hydrogen atoms bond together to form a hydrogen molecule, the two s-orbitals overlap to formation of a variant of a form a new molecular orbital. The two electrons then exist in this new orbital. The highest electron covalent bond, known as a density is between the two nuclei. ‘polar’ covalent bond (see Topic 3B.2). H H H H diagram showing orbital space filling model of a overlap hydrogen molecule ▲ fig C Formation of the σ bond in hydrogen. M03_IASL_CHEM_44860_TPC3_064-097.indd 72 01/06/2018 12:05  TOPIC 3 3B.1 COVALENT BONDING 73 EXAMPLE 2. CHLORINE A chlorine atom has an electronic configuration of 1s2 2s2 2p6 3s2 3px2 3py2 3pz1. When two chlorine atoms bond together, the two p orbitals (each containing a single electron) overlap. DID YOU KNOW? This view of the bonding in chlorine is very simple. An alternative theory Cl Cl Cl Cl describes the bonding as the overlap between two sp3 hybrid orbitals. This theory is not included in this book, but diagram showing the space filling model of a we would encourage you to orbital overlap chlorine molecule carry out your own research ▲ fig D Formation of the σ bond in chlorine. if you are interested. Look up ‘orbital hybridisation’. EXAMPLE 3. π BOND FORMATION Once a σ bond has been formed, it is possible, in certain circumstances, for a π bond to form. The π bond results in a high electron density both above and below the molecule, as shown in fig E. electron density above and sideways overlap of p-orbitals below the molecule ▲ fig E Formation of a π bond. This is what happens in the ethene molecule. One of the bonds between the carbon atoms is a σ bond; the other is a π bond. The π bond in ethene is weaker than the σ bond. This is the reason for the increased reactivity of alkenes compared with alkanes, and why alkenes can easily undergo addition reactions. (See Topic 5 for more information.) The triple bond in the nitrogen molecule (N ≡ N) is made up of one σ bond and two π bonds. p-orbitals p bond N N s bond p-orbital p-orbital The p-orbitals marked are those that are involved in the formation of the pi bonds. ▲ fig F Formation of the two π bonds in nitrogen. M03_IASL_CHEM_44860_TPC3_064-097.indd 73 01/06/2018 12:06  74 3B.1 COVALENT BONDING TOPIC 3 BOND LENGTH AND BOND STRENGTH The bond length is the distance between the nuclei of two atoms that are covalently bonded together. The strength of a covalent bond is measured in terms of the amount of energy required to break one mole of the bond in the gaseous state (see Topic 6). Table A shows the relationship between the bond length and bond strength of a selection of covalent bonds. BOND BOND LENGTH / nm BOND STRENGTH / kJ mol−1 LEARNING TIP Cl−Cl 0.199 242 When making a comparison between bond length and bond Br−Br 0.228 193 strength, it is important to I−I 0.267 151 compare ‘like with like’, in other words, compare things that are C−C 0.154 347 the same or very similar. C=C 0.134 612 For example, the strength of the C≡C 0.120 838 C—C bond (347 kJ mol−1) is greater than that of the N—N N−N 0.145 158 bond (158 kJ mol−1) despite N=N 0.120 410 being longer (0.154 nm compared with 0.145 nm). N≡N 0.110 945 In a molecule such as hydrazine O−O 0.148 144 (H2N—NH2), each nitrogen atom has a non-bonding (lone) pair O=O 0.121 498 of electrons and these repel one table A Relationship between bond length and bond strength for a range of covalent bonds. another, weakening the bond. In a molecule such as ethane The general relationship between bond length and bond strength, for bonds that are of a similar (H3C—CH3), the carbon atoms nature, is the shorter the bond, the greater the bond strength. This is a result of an increase in do not have any lone pairs. electrostatic attraction between the two nuclei and the electrons in the overlapping atomic orbitals. CHECKPOINT 1. Suggest a reason for the following trend in bond strengths: C—C > Si—Si > Ge—Ge 2. T  he F—F bond in fluorine is much shorter (0.142 nm) than the Cl—Cl (0.199 nm) bond in chlorine, and yet it is much weaker (158 kJ mol−1 compared with 242 kJ mol−1). Suggest a reason for this. 3. Suggest a reason why the sigma (σ) bond between the two carbon atoms in the ethene molecule is stronger that the pi (π) bond. SUBJECT VOCABULARY bond length the distance between the nuclei of two atoms that are covalently bonded together M03_IASL_CHEM_44860_TPC3_064-097.indd 74 01/06/2018 12:06  1 VELOCITY AND ACCELERATION 1A MOTION 75 SPECIFICATION 3B 2 ELECTRONEGATIVITY AND REFERENCE 3.10(ii) 3.13 3.14 BOND POLARITY LEARNING OBJECTIVES ◼ Know that electronegativity is the ability of an atom to attract a bonding pair of electrons. ◼ Know that ionic and covalent bonding are the extremes of a continuum of bonding type and that electronegativity differences lead to bond polarity. ◼ Understand what a polar covalent bond is. ◼ Understand that electron density maps for discrete (simple) molecules show that there is a high electron density between the nuclei of two covalently bonded atoms. WHAT IS ELECTRONEGATIVITY? Electronegativity is the ability of an atom to attract a bonding pair of electrons. The electronegativity of elements, in general: decreases down a group of the Periodic Table, that is, from top to bottom increases from left to right across a period. This is demonstrated in the following section of the Periodic Table (fig A). H He 2.1 Li Be B C N O F Ne 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Na Mg Al Si P S Cl Ar 0.9 1.2 1.5 1.8 2.1 2.5 3.0 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 0.8 1.0 1.3 1.5 1.6 1.6 1.5 1.8 1.8 1.8 1.9 1.6 1.6 1.8 2.0 2.4 2.8 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 0.8 1.0 1.2 1.4 1.6 1.8 1.9 2.2 2.2 2.2 1.9 1.7 1.7 1.8 1.9 2.1 2.5 Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 0.7 0.9 1.1 1.3 1.5 1.7 1.9 2.2 2.2 2.2 2.4 1.9 1.8 1.8 1.9 2.0 2.2 ▲ fig A able of electronegativities. DISTRIBUTION OF ELECTRON DENSITY If two atoms of the same element are bonded together by the overlap of atomic orbitals, the distribution of electron density between the two nuclei will be symmetrical. This is because the ability of each atom to attract the bonding pair of electrons is identical. The diagram in fig B is an electron density map for chlorine (Cl2): nuclei of chlorine atoms ▲ fig B Electron density map of a chlorine molecule. The diagram looks like a contour map. The contour lines correspond to electron density. You can think of them as showing how likely it is that a bonding electron will fall within that contour at a given instant in time. For a normal covalent bond, the contours are symmetrical around the nuclei. M03_IASL_CHEM_44860_TPC3_064-097.indd 75 01/06/2018 12:06  76 3B.2 ELECTRONEGATIVITY AND BOND POLARITY TOPIC 3 POLAR COVALENT BONDS However, if the two atoms bonded together are from elements that have different electronegativities, then the distribution of electron density will not be symmetrical about the two nuclei. This is shown in fig C by the electron density map for the hydrogen chloride (HCl) molecule. nucleus of nucleus of hydrogen atom chlorine atom (electronegativity 2.1) (electronegativity 3.0) ▲ fig C Electron density map of a hydrogen chloride molecule. The contour lines are more closely spaced near to the chlorine atom, which is the atom with the higher electronegativity. LEARNING TIP Since the electron density is higher around the chlorine atom, that end of the molecule has acquired a slightly negative charge. This is represented by the symbol δ–. The other end of the molecule Covalent bonds that are non- polar are sometimes called carries a slightly positive charge, represented by the symbol δ+. ‘normal’ covalent bonds or Hδ+ — Clδ− ‘pure’ covalent bonds. This is to distinguish them from polar A bond like this is called a polar covalent bond or sometimes just a ‘polar bond’. covalent bonds. Another way of representing a polar covalent bond is to use an arrow to show the direction of Both types of bond are formed electron drift. by the overlap of atomic orbitals. H Cl In both cases the overlapping Other examples of polar covalent bonds are: area contains two electrons per bond formed. Cδ+ Clδ−   Hδ+ Oδ−   Hδ+ Nδ−   Hδ+ Cδ– CONTINUUM OF BONDING TYPE Polar covalent bonds can be thought of as being between two ideals of bonding types. These ideals are: pure (100%) covalent pure (100%) ionic. Consider a polar covalent bond as a covalent bond that has some degree of ionic character. If the electronegativity difference is large enough, then the main type of bonding is ionic. A very approximate measure of the degree of ionic bonding in a compound is given in table A. ELECTRONEGATIVITY DIFFERENCE 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 1.1 APPROXIMATE % IONIC CHARACTER 0 0.5 1 2 4 6 9 12 15 19 22 26 ELECTRONEGATIVITY DIFFERENCE 1.2 1.3 1.4 1.5 1.6 1.7 1.8 1.9 2.0 2.1 2.2 2.3 APPROXIMATE % IONIC CHARACTER 30 34 39 43 47 51 55 59 63 67 70 74 ELECTRONEGATIVITY DIFFERENCE 2.4 2.5 2.6 2.7 2.8 2.9 3.0 3.1 3.2 3.3 APPROXIMATE % IONIC CHARACTER 76 79 82 84 86 88 89 90 91 92 table A Relationship between percentage ionic character and difference in electronegativity. M03_IASL_CHEM_44860_TPC3_064-097.indd 76 01/06/2018 12:06  TOPIC 3 3B.2 ELECTRONEGATIVITY AND BOND POLARITY 77 CHECKPOINT 1. Suggest why the electronegativity of fluorine is greater than that of chlorine, despite the fact that the nucleus of a chlorine atom contains more protons. 2. Ionic bonding and covalent bonding are two extremes of chemical bonding. Many compounds have bonding that is intermediate in character. (a) Giving an example in each case, explain what is meant by the terms: (i) ionic bonding, and (ii) covalent bonding. (b) Select a compound that has bonding of an intermediate character and explain why it has this type of bonding. 3. Place the following bonds in order of decreasing polarity (i.e. place the most polar first). C—Br  C—Cl  C—F  C—I Explain how you arrived at your answer. SUBJECT VOCABULARY electronegativity the ability of an atom to attract a bonding pair of electrons in a covalent bond polar covalent bond a type of covalent bond between two atoms where the bonding electrons are unequally distributed; because of this, one atom carries a slight negative charge and the other a slight positive charge M03_IASL_CHEM_44860_TPC3_064-097.indd 77 01/06/2018 12:06  SPECIFICATION 3B 3 BONDING IN DISCRETE REFERENCE 3.11(i) (SIMPLE) MOLECULES LEARNING OBJECTIVES ◼ Understand what is meant by the term discrete (simple) molecule. ◼ Draw dot-and-cross diagrams to show electrons in discrete molecules with single, double and triple bonds. ◼ Draw displayed formulae to represent the bonding in discrete molecules. DID YOU KNOW? DISCRETE MOLECULES Molecules are common in A discrete (simple) molecule is an electrically neutral group of two or more atoms held together organic substances (and by covalent bonds. therefore biochemistry). They also make up most of the DOT-AND-CROSS DIAGRAMS oceans and the atmosphere. Covalent and polar covalent bonding in discrete molecules can be shown by dot-and-cross diagrams. However, ionic crystals (salts) and giant covalent crystals Fig A shows the example of hydrogen, H2. (network solids) are often Further examples of dot-and-cross diagrams are shown in table A. made up of repeating unit cells that extend either in a SUBSTANCE DOT-AND-CROSS DIAGRAM plane (such as in graphene) Water, H2O H O H or three-dimensionally (such as in diamond or H O H sodium chloride). These substances are not composed of molecules. Solid metals are also not composed of Ammonia, NH3 H molecules. H N H H nucleus of H H hydrogen atom (electronegativity 2.1) H N H ▲ fig A Dot-and-cross diagram for hydrogen with overlapping circles. Methane, CH4 H H C H H H H C H H table A Dot-and-cross diagrams for water, ammonia and methane. EXAM HINT When drawing a molecule such as chloromethane, do not forget to show all of the non-bonding electrons on the chlorine atom. M03_IASL_CHEM_44860_TPC3_064-097.indd 78 01/06/2018 12:06  TOPIC 3 3B.3 BONDING IN DISCRETE (SIMPLE) MOLECULES 79 THE OCTET RULE You might read that in order to form a stable compound, the outer shell of each atom must have the same number of electrons as the outer shell of a noble gas. In most cases this will be eight electrons. This has led to a rule that is often referred to as the ‘octet rule’. This is not always true, as you can see from the examples in table B. In each case, the outer shell of the central atom of the molecule does not contain eight electrons. SUBSTANCE DOT-AND-CROSS DIAGRAM NUMBER OF ELECTRONS AROUND CENTRAL ATOM Beryllium chloride, BeCl2 Cl Be Cl 4 Boron trichloride, BCl3 Cl 6 Cl B Cl Phosphorus(V) chloride, PCl5 10 Cl Cl Cl P Cl Cl Sulfur hexafluoride, SF6 F 12 S F F F F F table B Examples breaking the octet rule. DOT-AND-CROSS DIAGRAMS OF MOLECULES CONTAINING DISPLAYED FORMULAE (FULL STRUCTURAL FORMULAE) MULTIPLE BONDS A displayed (full structural) formula shows each bonding pair as a line drawn between the two atoms involved. Table C shows the dot-and-cross diagrams for three molecules (O2, N2, CO2) that contain a double or triple bond. Table C gives some examples of dot-and-cross diagrams together with the displayed formulae. SUBSTANCE DOT-AND-CROSS DIAGRAM DISPLAYED FORMULA Water, H2O H O H H—O—H LEARNING TIP Ammonia, NH3 H H Although it is essential to show H N H all of the non-bonding (lone) H—N—H pairs of electrons in a dot-and- Oxygen, O2 O O O=O cross diagram, it is not necessary to show them in a Nitrogen, N2 N≡N displayed formula. N N SUBJECT VOCABULARY Carbon dioxide, CO2 O C O O=C=O discrete (simple) molecule an table C Examples of displayed formulae with the corresponding dot-and-cross diagram. electrically neutral group of two or more atoms held together by covalent bonds CHECKPOINT displayed (full structural) 1. Draw a dot-and-cross diagram for each of the following molecules: formula a formula that shows each bonding pair as a line (a) H2S  (b) PH3  (c) PF3  (d) SCl2  (e) AsF5  (f) HCN  (g) SO2 drawn between the two atoms 2. Draw the displayed formula for each of the molecules in Question 1. involved M03_IASL_CHEM_44860_TPC3_064-097.indd 79 01/06/2018 12:06  SPECIFICATION REFERENCE 3B 4 DATIVE COVALENT BONDS 3.11(ii) LEARNING OBJECTIVES ◼ Know that in a dative covalent bond both electrons in the bond are supplied by only one of the atoms involved in forming the bond. ◼ Be able to draw dot-and-cross diagrams for some molecules and ions that contain dative covalent bonds, including Al2Cl6 and the ammonium ion. DATIVE COVALENT BOND FORMATION A dative covalent bond is formed when an empty orbital of one atom overlaps with an orbital containing a non-bonding pair (lone pair) of electrons of another atom. The bond is often represented by an arrow from the atom providing the pair of electrons, to the atom with the empty orbital. Below are three examples of dative covalent bonds. THE HYDROXONIUM ION, H3O+ The dot-and-cross diagram and the displayed formula of a hydroxonium ion are shown in fig A. H O H 1 H O H 1 H H ▲ fig A Dot-and-cross diagram and displayed formula for the hydroxonium ion. The empty 1s orbital of the H+ ion overlaps with the orbital of the oxygen atom that contains the lone pair of electrons. THE AMMONIUM ION, NH4+ The dot-and-cross diagram and the displayed formula of an ammonium ion are shown in fig B. H 1 H 1 H N H H N H H H ▲ fig B Dot-and-cross diagram and displayed formula for the ammonium ion. The empty 1s orbital of the H+ ion overlaps with the orbital of the nitrogen atom that contains the lone pair of electrons. ALUMINIUM CHLORIDE, Al2Cl6 The aluminium atom in the AlCl3 molecule has only six electrons in its outer shell and so has an empty orbital (fig C). Cl Cl Al Cl ▲ fig C Dot-and-cross diagram for aluminium chloride. In the gas phase, just above its sublimation temperature, aluminium chloride exists as Al2Cl6 molecules (fig D). M03_IASL_CHEM_44860_TPC3_064-097.indd 80 01/06/2018 12:06  TOPIC 3 3B.4 DATIVE COVALENT BONDS 81 Two AlCl3 molecules bond together. One of the atomic orbitals of a chlorine atom of one AlCl3 molecule that contains a lone pair overlaps with the empty orbital of the aluminium atom of a second AlCl3 molecule. The same happens between the chlorine atom of the second molecule and the aluminium atom of the first molecule. One chlorine atom from each molecule acts as a bridge connecting the two molecules with dative covalent bonds. Cl Cl Cl Al Al Cl Cl Cl ▲ fig D Displayed formula for the aluminium dimer. CHECKPOINT 1. (a) Draw a dot-and-cross diagram for a molecule of NH3 and a molecule of BF3. (b) Draw a dot-and-cross diagram for a molecule of NH3·BF3.  raw a dot-and-cross diagram and displayed formula for the AlCl4− ion and identify the dative 2. D covalent bond. 3. O  ne way of describing the bonding in a molecule of carbon monoxide (CO) is to state that it contains two covalent bonds and one dative bond. Using this description, draw a dot-and-cross diagram and displayed formula for a molecule of carbon monoxide. SUBJECT VOCABULARY dative covalent bond the bond formed when an empty orbital of one atom overlaps with an orbital containing a lone pair of electrons of another atom M03_IASL_CHEM_44860_TPC3_064-097.indd 81 01/06/2018 12:06  SPECIFICATION 3C REFERENCE 1 SHAPES OF MOLECULES 3.16 3.17 PART 3.18 3.19 AND IONS If each double bond is treated as an electron pair, then the LEARNING OBJECTIVES molecule is linear, like BeCl2. ◼ Understand the principles of the electron-pair NUMBER OF NUMBER OF SHAPE EXAMPLE repulsion theory, used to interpret and predict the BOND PAIRS LONE PAIRS shapes of simple molecules and ions. 2 0 linear Cl Be Cl ◼ Understand the term bond angle. Cl ◼ Know and be able to explain the shapes of, and bond trigonal 3 0 B angles in, BeCl2, BCl3, CH4, NH3, NH4+, H2O, CO2, planar gaseous PCl5, SF6 and C2H4. Cl Cl ◼ Be able to apply the electron-pair repulsion theory to H predict the shapes of, and bond angles in, molecules 4 0 tetrahedral C H and ions analogous to those mentioned above. H H ELECTRON PAIR REPULSION THEORY Cl The electron pair repulsion (EPR) theory states that: Cl trigonal P the shape of a molecule or ion is caused by repulsion between 5 0 Cl bipyramidal Cl the pairs of electrons, both bond pairs and lone (non-bonding) Cl pairs, that surround the central atom the electron pairs arrange themselves around the central atom F so that the repulsion between them is at a minimum F F 6 0 octahedral S lone pair–lone pair repulsion > lone pair–bond pair repulsion > bond pair–bond pair repulsion. F F F LEARNING TIP N This theory is sometimes also called the valence shell electron pair trigonal 3 1 H H repulsion theory, abbreviated to VSEPR. pyramidal H The first two rules are used to obtain the basic shape of the O molecule or ion. The third rule is used to estimate values for the 2 2 V-shaped bond angles. H H table A Shapes of molecules. THE SHAPES OF MOLECULES AND IONS To obtain the shape of a molecule or ion it is first necessary to EXAMPLE 2. ETHENE, C2H4 obtain the number of bond pairs and lone pairs of electrons The displayed formula of ethene is: around the central atom. H H The easiest way to do this is by drawing a dot-and-cross diagram. C C You can then apply the guidelines listed in table A. H H MOLECULES WITH MULTIPLE BONDS There are no lone pairs on either carbon atom. To determine the shape of a molecule containing one or more multiple bonds, treat each multiple bond as if it contained only Treating each double bond as an electron pair produces a planar one pair of electrons. molecule with 120° bond angles. H H EXAMPLE 1. CARBON DIOXIDE, CO2 C C The displayed formula for carbon dioxide is O=C=O. There are no lone pairs on the carbon atom. H H M03_IASL_CHEM_44860_TPC3_064-097.indd 82 01/06/2018 12:06  TOPIC 3 3C.1 SHAPES OF MOLECULES AND IONS 83 THE BOND ANGLES IN MOLECULES AND IONS Table B shows the bond angles of a range of molecules and ions. Linear, e.g. BeCl2 180° Trigonal planar, e.g. BCl3 Cl The bond angle is 180°. The bond angle is 120°. 120° Cl Be Cl B Cl Cl Tetrahedral, e.g. CH4 H Trigonal pyramidal, e.g. NH3 N The bond angle is 109.5°. 109.5° The bond angle is 107°. H H C H H 107° H H Lone pair–bond pair repulsion is greater than bond pair–bond pair repulsion, so the angle is slightly less than 109.5°. V-shaped, e.g. H2O Trigonal bipyramidal, e.g. PCl5 The bond angle is 104.5°. There are two bond angles: 90° and 120°. O Cl 90° Cl H 104.5° H Cl P 120° Lone pair–lone pair repulsion is greater than Cl lone pair–bond pair repulsion, so the bond angle Cl is even further depressed from 109.5°, and is slightly less than the 107° in NH3. Octahedral, e.g. SF6 Tetrahedral, e.g. NH4+ There are two bond angles: 90° and 180°. As with CH4, the bond angles are 109.5°. F H 1 90° F F S H N F F F H 109.5° H The angle between the bonds of two fluorine Note the change from 107° in ammonia to atoms opposite one another is 180°. 109.5° in the ammonium ion. table B The bond angles of a range of molecules and ions. CHECKPOINT 1. (a) Draw a diagram to show the shape of each of the following molecules: (i) H2S  (ii) PH3  (iii) PF3  (iv) SCl2  (v) AsF5  (vi) HCN  (vii) SO2 (b) Give the name of each shape. SUBJECT VOCABULARY 2. Solid phosphorus pentachloride has the formula [PCl4]+(PCl6]−. electron pair repulsion (EPR) (a) Draw a diagram to show the shape of each ion. theory the electron pairs on (b) State the bond angles present in each ion. the central atom of a molecule or ion arrange themselves in 3. Two possible ways of arranging the bonding pairs and lone pairs of electrons in a molecule of XeF4 are: order to create the minimum repulsion between them; lone F F F pair-lone pair repulsion is Xe and Xe greater than lone pair-bond F F F F repulsion, which in turn is F greater than bond pair-bond Suggest which of these two arrangements is the more likely and justify your answer. pair repulsion M03_IASL_CHEM_44860_TPC3_064-097.indd 83 01/06/2018 12:06  SPECIFICATION REFERENCE 3C 2 NON-POLAR AND POLAR MOLECULES 3.15 PART POLYATOMIC MOLECULES LEARNING OBJECTIVES 1. LINEAR MOLECULES ◼ Understand the difference between non-polar and Example: carbon dioxide, CO2 polar molecules and be able to predict whether or not Both bonds in the carbon dioxide molecule are polar, but the a given molecule is likely to be polar. dipoles cancel out one another. Oδ−=Cδ+=Oδ− SHAPE AND POLARITY The drift of bonded electrons towards the more electronegative ▲ fig A Dipoles in carbon dioxide. element (see Topic 3B.2) results in a separation of charge. This separation of charge is called a dipole. The carbon dioxide molecule is therefore non-polar. Each of the bonds in a molecule has its own dipole associated 2. TRIGONAL PLANAR MOLECULES with it. The overall dipole of a molecule depends on its shape. Example: boron chloride, BCl3 Depending on the relative angles between the bonds, the individual dipoles can either reinforce one another or cancel out Clδ2 each other. B δ1 If the cancellation is complete, the resulting molecule will have Cl δ2 Clδ2 no overall dipole and is said to be ‘non-polar’. If the dipoles reinforce one another, the molecule will possess ▲ fig B Dipoles in boron trichloride. an overall dipole and is said to be ‘polar’. All three B—Cl bonds are polar, but because the molecule is symmetrical the dipoles cancel out one another. The molecule is DIATOMIC MOLECULES non-polar. Hydrogen and chlorine are examples of diatomic molecules that are non-polar. The two atoms in each molecule are the same and 3. TETRAHEDRAL MOLECULES so have the same electronegativity. The distribution of electron Example 1: tetrachloromethane, CCl4 density of the bonding electrons in either molecule is totally Clδ2 symmetrical (see Topic 3B.2). The bond in each is therefore non- polar, making the molecules non-polar. δ1 C Clδ2 However, the bond in the hydrogen chloride molecule is polar Clδ2 Clδ2 because the electronegativity of chlorine (3.0) is greater than that of hydrogen (2.1). ▲ fig C Dipoles in tetrachloromethane. Hδ+ Clδ− All four C—Cl bonds are polar, but because the molecule is symmetrical the dipoles cancel out one another. The molecule is Since this is the only polar bond in the molecule, the molecule non-polar. itself is polar. Example 2: trichloromethane, CHCl3 The following symbol is used to represent a dipole: H The dipole in the hydrogen chloride molecule is shown as: C Hδ+ — Clδ− Cl Cl Cl ▲ fig D Dipoles in trichloromethane. All four bonds are polar but, although the molecule is symmetrical, the dipoles reinforce one another and so the molecule is polar. M03_IASL_CHEM_44860_TPC3_064-097.indd 84 01/06/2018 12:06  TOPIC 3 3C.2 NON-POLAR AND POLAR MOLECULES 85 4. TRIGONAL PYRAMIDAL MOLECULES CHECKPOINT Example: ammonia, NH3 1. A bond between two atoms in a molecule may possess a dipole. All three N—H bonds are polar and the dipoles reinforce one (a) Explain how this dipole arises. another. The molecule is polar. (b) Some bonds that you are likely to meet in organic chemistry N are listed. Which of these bonds are likely to possess a dipole? H In each case indicate which atom is δ+ and which is δ−. H H C—Cl O—H C—C C—O C=C C—N N—H ▲ fig E Dipoles in ammonia. 2. State whether each of the following molecules are non-polar or polar. In each case, explain your reasoning. 5. V-SHAPED MOLECULES (a) H2S Example: water, H2O (b) CH4 O (c) SO2 H H (d) SO3 ▲ fig F Dipoles in water. (e) AlBr3 (f) PBr3 Both O—H bonds are polar and the dipoles reinforce one another. The molecu