Unit 4 Chemical Bonding and Structure PDF

04 Chemical bonding and structure Essential ideas Scanning electron micrograph of graphite shows its layered 4.1 Ionic compou...

04 Chemical bonding and structure Essential ideas Scanning electron micrograph of graphite shows its layered 4.1 Ionic compounds consist of ions held together in lattice structures by structure. Graphite is a form ionic bonds. of pure carbon, in which the bonding of atoms within a 4.2 Covalent compounds form by the sharing of electrons. layer is much stronger than the forces between the layers. An understanding of bonding 4.3 Lewis (electron dot) structures show the electron domains in the enables scientists to explain valence shell and are used to predict molecular shape. and predict many of the properties of materials. 4.4 The physical properties of molecular substances result from different types of forces between their molecules. 4.5 Metallic bonds involve a lattice of cations with delocalized electrons. 14.1 Larger structures and more in-depth explanations of bonding systems often require more sophisticated concepts and theories of bonding. 14.2 Hybridization results from the mixing of atomic orbitals to form the same number of new equivalent hybrid orbitals that can have the same mean energy as the contributing atomic orbitals. We learned in Chapter 2 that all elements are made of atoms but that there are only about 100 chemically different types of atom. Yet we know that we live in a world made up of literally millions of different substances: somehow these must all be formed from just these 100 atomic building blocks. The extraordinary variety arises from the fact that atoms readily combine with each other and they do so in a myriad of different ways. They come together in small numbers or large, with similar atoms or very different atoms, but always the result of the combination is a stable association known as a chemical bond. Atoms linked together by bonds therefore have very different properties from their parent atoms. In this chapter we will study the main types of chemical bonds – the ionic bond, the covalent bond, and the metallic bond – and also consider other forces that help to hold substances A molecule of insulin, the together. Our study of the covalent bond at this level will use some of the concepts hormone essential for the regulation of glucose in the from quantum mechanical theory developed in Chapter 2 to explain the shapes and body. The ball-and-stick properties of molecules in more detail. As electrons are the key to the formation of all model shows all the atoms these bonds, a solid understanding of electron configurations will help you. and bonds within the protein molecule. Insulin was the Chemical reactions take place when some bonds break and others re-form. Being first protein to have its entire able to predict and understand the nature of the bonds within a substance is therefore structure elucidated. central to explaining its chemical reactivity. 139  04 Chemical bonding and structure 4.1 Ionic bonding and structure Understandings: Positive ions (cations) form by metals losing valence electrons. Negative ions (anions) form by non-metals gaining electrons. The number of electrons lost or gained is determined by the electron configuration of the atom. The ionic bond is due to electrostatic attraction between oppositely charged ions. Under normal conditions, ionic compounds are usually solids with lattice structures. Applications and skills: Deduction of the formula and name of an ionic compound from its component ions, including polyatomic ions. Guidance Students should be familiar with the names of these polyatomic ions: NH4+, OH–, NO3–, HCO3–, CO32–, SO42–, and PO43– Explanation of the physical properties of ionic compounds (volatility, electrical conductivity, and solubility) in terms of their structure. Ions form when electrons are transferred An ion is a charged All atoms are electrically neutral, even though they contain charged particles known particle. Ions form from as protons and electrons. This is because the number of protons (+) is equal to the atoms or from groups of number of electrons (−), and so their charges cancel each other out. The positively atoms by loss or gain of charged protons, located within the nucleus of the atom, are not transferred during one or more electrons. chemical reactions. Electrons, however, positioned outside the nucleus, are less tightly held and outer electrons, known as valence electrons, can be transferred when atoms react together. When this happens the atom is no longer neutral, but instead carries an electric charge and is called an ion. The charge on the ion which forms is therefore determined by how many electrons are lost or gained. We learned in Chapter 3 that the group number in the Periodic Table relates to the number of electrons in the outer shell of the atoms of all the elements in that group. We also learned that Group 18 elements, known as the noble gases, where the atoms all have full outer shells of electrons, are especially stable and have almost no tendency It may help you to remember: CATion is to react at all. This full outer shell behaves in a sense like the ‘ultimate goal’ for other PUSSYtive and aNion is atoms: they react to gain the stability associated with this by losing or gaining the Negative. appropriate number of electrons, whichever is the easiest (in energetic terms). An energy-efficient light bulb. A compact fluorescent lamp has a gas-filled glass tube which contains a small amount of mercury vapour mixed with argon under low pressure. As argon is a noble gas, it is very unreactive owing to its stable electron arrangement, so it helps to provide an inert environment. 140  Elements that have a small number of electrons in their outer shells (Groups 1, 2, and When an atom loses 13) will lose those electrons and form positive ions called cations. These elements are electrons it forms a the metals. positive ion, called a cation. When an atom Elements that have higher numbers of electrons in their outer shells (Groups 15, 16, gains electrons it forms and 17) will gain electrons and form negative ions called anions. These elements are a negative ion, called the non-metals. an anion. The number We are now able to summarize how the position of an element in the Periodic Table of charges on the ion formed is equal to the enables us to predict the type of ion that it will form. number of electrons lost or gained. Group Example Number Electrons Number of Charge Type of number of valence lost or electrons on ion element electrons gained transferred formed 1 sodium 1 lost 1 1+ metal 2 calcium 2 lost 2 2+ metal 13 aluminium 3 lost 3 3+ metal 14 carbon 4 – – – non-metal 15 phosphorus 5 gained 3 3– non-metal 16 oxygen 6 gained 2 2– non-metal Metals form cations by losing valence electrons. 17 bromine 7 gained 1 1– non-metal Non-metals form anions by gaining electrons. Note that elements in Group 14, having four electrons in their outer shell, do not have a tendency to gain or to lose electrons, and so they generally do not form ions. This is Do we have direct or because the energy involved in transferring four electrons would simply be too large to indirect evidence for the existence of ions? Is there be favourable. These elements therefore react to form a different type of bond, which an important difference we will discuss later in this chapter. between the two types of evidence? Worked example Refer to the Periodic Table to deduce the charge on the ion formed when the following elements react: (i) lithium (ii) sulfur (iii) argon Solution (i) lithium is in Group 1 so forms Li+ (ii) sulfur is in Group 16 so forms S2– (iii) argon is in Group 18 so does not form ions For some elements though, it is difficult to predict the ion that will form from its position in the Periodic Table. For example, as we learned in Chapter 3, the metals occurring in the middle of the Periodic Table, known as the transition elements, have an electron configuration that allows them to lose different numbers of electrons from their d sub-shell and so form stable ions with different charges. The transition element iron, Fe, for example, can form Fe2+ by losing two electrons or Fe3+ by losing three electrons, depending on the reacting conditions. The two ions have distinct properties, such as forming compounds with different colours. 141  04 Chemical bonding and structure Compounds containing different ions of iron can be distinguished by colour: the left beaker contains Fe2+(aq) and the right beaker Fe3+(aq). Similar colour changes occur when iron rusts as it reacts with oxygen to form these different ions. Likewise, the element copper can exist as Cu2+ and Cu+ and again these ions can be distinguished by colour. Other examples of elements that form ions that are not obvious from their group number are: lead, Pb, despite being in Group 14, forms a stable ion Pb2+ tin, Sn, also in Group 14, can form Sn4+ and Sn2+ Fehling’s reagent uses the silver, Ag, forms the ion Ag+ different colours of the copper ions to test for simple sugars. hydrogen, H, can form H– (hydride) as well as the more common H+. The left tube containing the When the charge on an ion needs to be specified, the oxidation number is given in blue Cu2+ ion changes to the red Cu+ ion seen on the right Roman numerals in brackets after the name of the element. For example, the red Cu+ when warmed with glucose or ion shown above can be written as copper(I) oxide. Oxidation numbers are explained other ‘reducing sugars’. fully in Chapter 9. Finally, there are some ions that are made up of more than one atom which together have experienced a loss or gain of electrons and so carry a charge. These species When writing the symbol are called polyatomic ions, and many of them are found in commonly occurring for ions, note the charge is written as a superscript compounds. It will help you to become familiar with the examples in the table below, with the number first and as you will often use them when writing formulas and equations. (Note that this the charge next, e.g. N3–. information is not supplied in the IB data booklet.) When an ion X carries a charge of 1+ or 1– it is Polyatomic ion Charge Symbol Example of compound containing written just as X+ or X–. name on ion this ion nitrate 1– NO3– lead nitrate hydroxide 1– OH– barium hydroxide hydrogencarbonate 1– HCO3– potassium hydrogencarbonate Note the common names carbonate 2– CO32– magnesium carbonate of some compounds sulfate 2– SO42– copper sulfate give a clue to their composition. Here you phosphate 3– PO43– calcium phosphate can see that the ending ammonium 1+ NH4+ ammonium chloride ‘-ate’ refers to ions that contain oxygen bonded to another element. We will learn how to write the formulas for these compounds in the next section. 142 Ionic compounds form when oppositely charged ions attract Ions do not form in isolation. Rather the process of ionization – where electrons are transferred between atoms – occurs when an atom that loses electrons passes them directly to an atom that gains them. Typically, this means electrons are transferred from a metal element to a non-metal element. For example, sodium (metal) passes an electron to chlorine (non-metal) when they react together. " Na ! Cl [Na]! [ Cl ] 1s22s22p63s1 1s22s22p63s23p5 1s22s22p6 1s22s22p63s23p6 sodium chlorine sodium chloride atom atom ion ion Note that both ion products, Na+ and Cl–, have the electron configuration of a noble Coloured scanning electron gas. micrograph of crystals of table salt, sodium chloride, Now the oppositely charged ions resulting from this electron transfer are attracted to NaCl. The very reactive each other and are held together by electrostatic forces. These forces are known as an elements sodium and chlorine have combined together to ionic bond, and ions held together in this way are known as ionic compounds. form this stable compound Remember that in forming the ionic compound there is no net loss or gain of containing Na+ and Cl– ions. electrons, and so the ionic compound, like the atoms that formed it, must be electrically neutral. Writing the formula for the ionic compound therefore involves The ionic bond is due to balancing the total number of positive and negative charges, taking into account the different electrostatic attraction charges on each ion. between oppositely charged ions. For example, magnesium oxide is made up of magnesium ions Mg2+ and oxide ions O2–. Here each magnesium atom has transferred two electrons to each oxygen atom Note that in many non- and so the compound contains equal numbers of each ion. Its formula, Mg2+O2–, is metal elements the ending usually written as MgO. But when magnesium reacts with fluorine, each Mg loses two of the name changes electrons, whereas each F gains only one electron. So it will take two fluorine atoms to to ‘-ide’ when ions are combine with each magnesium atom, and the compound that results will therefore present. For example, chlorine (the element) have the ratio Mg : F = 1 : 2. This is written as Mg2+F–2 or MgF2. becomes chloride (the ion), oxygen becomes oxide, Worked example nitrogen becomes nitride, Write the formula for the compound that forms between aluminium and oxygen. etc. Solution Note the convention in 1 Check the Periodic Table for the ions that each element will form. naming ionic compounds aluminium in Group 13 will form Al3+; oxygen in Group 16 will form O2– that the positive ion 2 Write the number of the charge above the ion: 3 2 is written first and the negative ion second. Al O Cross-multiply these numbers 3 2 Al O Note that the formula of the compound shows the Or you can directly balance the charges: here you need 6 of each charge: simplest ratio of the ions it contains. So, for example, 2 × Al3+ = 6+ and 3 × O2– = 6– magnesium oxide is not 3 Write the final formula using subscripts to show the number of ions: Al2O3 Mg2O2 but MgO. 143  04 Chemical bonding and structure It is common practice to leave the charges out when showing the final formula. If the formula contains more than one polyatomic ion, brackets are used around the ion before the subscript. Worked example Write the formula for ammonium phosphate. Solution The compound contains two polyatomic ions, and you need to know these: NH4+ and PO43– Balancing the charges: 3 × NH4+ = 3+ and 1 × PO43– = 3– or: Following the steps above: 1 3 NH4+ PO43– So the formula is (NH4)3PO4. Exercises 1 Write the formula for each of the compounds in the table on page 142. 2 Write the formula for each of the following compounds: (a) potassium bromide (d) copper(II) bromide (b) zinc oxide (e) chromium(III) sulfate (c) sodium sulfate (f) aluminium hydride 3 Name the following compounds: (a) Sn3(PO4)2 (d) BaSO4 (b) Ti(SO4)2 (e) Hg2S (c) Mn(HCO3)2 4 What are the charges on the positive ions in each of the compounds in Q3 above? 5 What is the formula of the compound that forms from element A in Group 2 and element B in Group 15? 6 Explain what happens to the electron configurations of Mg and Br when they react to form the compound magnesium bromide. Ionic compounds have a lattice structure Cl! ion The forces of electrostatic attraction between ions in a compound cause the ions to surround themselves with ions of opposite charge. As a result, the ionic compound takes on a predictable three-dimensional crystalline structure known as an ionic Na" ion lattice. The details of the lattice’s geometry vary in different compounds, depending mainly on the sizes of the ions, but it always involves this fixed arrangement of ions Figure 4.1 The NaCl lattice is built up from oppositely based on a repeating unit. The term coordination number is used to express the charged sodium and chloride number of ions that surround a given ion in the lattice. For example, in the sodium ions. chloride lattice, the coordination number is six because each Na+ ion is surrounded by six Cl– ions and each Cl– ion is surrounded by six Na+ ions. Make sure that you avoid the term ‘molecular Note that the lattice consists of a very large number of ions and it can grow indefinitely. formula’ when describing As ionic compounds do not therefore exist as units with a fixed number of ions, their ionic compounds, but formulas are simply an expression of the ratio of ions present. The simplest ratio is known instead use the term formula unit. as the formula unit, which is an empirical formula, as described in Chapter 1. 144 We will learn in Chapter 5 that lattice energy is a measure of the strength of attraction between the ions within the lattice. This is Computer graphic of greater for ions that are small and crystallized common salt, highly charged, as they have a larger NaCl. Small spheres represent charge density. Na+ ions and larger spheres Cl–. The lattice is arranged so each Na+ ion has six oppositely The physical properties charged nearest neighbours and vice versa. of ionic compounds reflect their lattice structure Physical properties are those that can be examined without chemically altering the substance. Our knowledge of ionic bonds and lattice structure helps us to interpret and explain some of these properties for ionic compounds. Melting points and boiling points Ionic liquids are efficient Ionic compounds tend to have high melting and boiling points as the forces of solvents with low volatility electrostatic attraction between the ions in the lattice are strong and so require large and are being used amounts of heat energy to break. These compounds are therefore solids at room increasingly as solvents in Green Chemistry for temperature, and will only melt at very high temperatures. Sodium chloride, for energy applications and example, remains solid until about 800 °C. industrial processes. They are usually made of The melting and boiling points are generally higher when the charge on the ions is organic salts. greater, due to the increased attraction between the ions. For example, the table below compares the melting points of sodium oxide and magnesium oxide. Ionic compound Charge on metal ion Melting point CHALLENGE Na2O 1+ 1132 °C YOURSELF MgO 2+ 2800 °C 1 The melting point of aluminium oxide, Al2O3, is The high melting points of ionic compounds become an economic consideration slightly lower than that of in many industrial processes, such as the electrolysis of molten ionic compounds magnesium oxide, despite the 3+ charge on the discussed in Chapter 9. It can be very expensive to maintain such high temperatures, aluminium ion. What factors and this is an important factor in considering suitable methods to extract a reactive might help to explain this? metal such as aluminium from its ore. Volatility is a term used to describe the tendency of a substance to vaporize. In Ionic compounds have summary, ionic compounds can be described as having a low volatility or being non- low volatility. volatile. Our everyday encounter with them as crystalline solids with low odour is consistent with this. Solubility Solubility refers to the ease with which a solid (the solute) becomes dispersed through a liquid (the solvent) to form a solution. You probably know from common Ionic compounds are observations that common salt, sodium chloride, readily dissolves in water but does generally soluble in not dissolve in oil. Why is this? There are several factors involved, but in general, ionic or polar solvents but not soluble in non- solubility is determined by the degree to which the separated particles of solute are polar solvents. able to form bonds or attractive forces with the solvent. 145  04 Chemical bonding and structure Consider an ionic compound being placed in water. As we will learn in the next Figure 4.2 Dissolving of section, water molecules are polar, which means they have some separation of charge NaCl in water involves the in their structure. At the contact surface, the partial charges in the water molecules are attraction of the polar water attracted to ions of opposite charge in the lattice, which may cause the ions to dislodge molecules to the opposite charged ions in the NaCl from their position. As these ions separate from the lattice, they become surrounded lattice, and the hydration of by water molecules and are said to be hydrated. When this happens the solid is the separated ions. dissolved. State symbols are used to show this change as follows: H H +H2O O NaCl(s) → NaCl(aq) H H +H2O O H O H NaCl(s) → Na+(aq) + Cl–(aq) Na+ If a liquid other than water is able to dissolve the solid, H the ions are said to be solvated and an appropriate state O H O symbol to denote the solvent is used. H H In the case of solvents like oil or hexane, C6H14, which H O Na+ Cl– Na+ H H are non-polar and so have no charge separation, there O H is no attraction between the liquid and the ions. So Cl– Na+ Cl– here the ions remain tightly bound to each other in the Na+ Cl– Na+ lattice, and the solid is insoluble. O H H H Cl– Na+ Cl– O This suggests that solubility trends are based on the H O H H similar chemical nature of the solute and solvent, as O H this is most likely to lead to successful interactions H Cl– between them. The expression ‘like dissolves like’ is often H H used to capture this notion. Note though that this is a O H H O generalized statement and somewhat over-simplified as there are some important exceptions. H H O Electrical conductivity The ability of a compound Condom conductivity test. to conduct electricity Condoms are tested for holes depends on whether it by being filled with water contains ions that are able and placed in a solution of to move and carry a charge. NaCl, and then attached to Ionic compounds are not electrodes. The current will not be conducted across able to conduct electricity the insulating material of the in the solid state as the ions condom, but if there is a hole are firmly held within the the current will be conducted lattice and so cannot move. into the salty water, triggering an alarm. All condoms are However, when the ionic conductivity tested in this way. compound is either present in the liquid state (molten), Ionic compounds do or dissolved in water not conduct electricity (aqueous solution), the ions in the solid state, but will be able to move. Therefore ionic compounds as liquids or aqueous solutions do conduct when molten or in aqueous solution. show electrical conductivity. 146 Brittleness NATURE OF SCIENCE Ionic compounds are usually brittle, which means the crystal tends to shatter when The theory that ionic force is applied. This is because movement of the ions within the lattice places ions of compounds have a lattice the same charge alongside each other, so the repulsive forces cause it to split. structure of discrete ions leads to some predictions about their physical Different ionic compounds have a different extent of properties. The fact that these predictions can be ionic character tested by observation A binary compound is one that contains only two elements. We have seen that in order and experiment is central to the scientific process, for any two elements to react to form an ionic compound, they must have very different and helps develop the tendencies to lose or gain electrons. A relatively simple assessment of this can be made by theory. In general, data looking at their positions on the Periodic Table. Metals on the left lose electrons most on the volatility, solubility, easily, while non-metals on the right gain electrons most easily. Also, we learned in and conductivity of ionic compounds as discussed Chapter 3 that the tendency to lose electrons and form positive ions increases down a here support the lattice group, whereas the tendency to gain electrons and form negative ions increases up a model. Yet there are some group. So the highest tendency to react together and form ionic compounds will be exceptions too, which between metals on the bottom left and non-metals on the top right of the Periodic Table. need to be explained. The fact that aluminium oxide’s increasing tendency to melting point is not as form negative ions high as would be expected most reactive non-metal from a model based purely 1 2 13 14 15 16 17 18 on Al3+ and O2– ions opens the question of what other non-metals types of bonding might be metals present, and why. Figure 4.3 The pairs of elements that react most easily increasing tendency to to form ionic compounds form positive ions are metals on the bottom most reactive metal left of the Periodic Table and non-metals on the top right, Another way to judge the tendency of two elements to form an ionic compound is indicated here by asterisks. by looking at electronegativity values. As explained in Chapter 3 on page 107, electronegativity is a measure of the ability of an atom to attract electrons in a covalent bond, and is described using the Pauling scale of values. These are given in section 8 of Figure 4.4 Periodic trends in electronegativity values show the IB data booklet and are summarized here. an increase along a period and 1 up a group. H 2.2 3 4 1.8 electronegativity difference