Chemical Bonding: Chemical Formulas PDF

05 Chemical Bonding: Chemical Formulas 5.1 Introduction Most of the matertatsaround us are made from combinations of elements, e.g. water is formed when hydrogen and oxygen combtne together, table salt is formed by co...

05 Chemical Bonding: Chemical Formulas 5.1 Introduction Most of the matertatsaround us are made from combinations of elements, e.g. water is formed when hydrogen and oxygen combtne together, table salt is formed by combtntng sodium and chlorine and carbon dioxide is formed when carbon and oxygen combine together. Substances made by joining together two or more different elements are called ilJir till! n compounds. Fig. 5.1 Helium and argon are two of the noble (inert) gases. Helium is a very light gas and is safe for use in filling balloons as it is so unreactive.Argon is used in medical lasers to correct A compound is a substance that is made up of two eye defects. It is also used to provide an inert atmosphere or more different elementscombined together for making silicon chips for the electronics industry. Light chemically. bulbs that contain filaments are usually filled with argon. If the bulb contained oxygen the filament would burn away. When elements combine to form compounds, there All of the noble gases consist of single atoms. This is are attractive forces that hold the atoms together in explained by saying that these elements do not react this new substance. These attractive forces are called because they have stable electron configurations. chemical bonds. In this chapter we will study the Since all of the noble gases (except helium) have eight different types of chemical bonds that exist. We will electrons in their outer energy level, the rule was called see how a knowledge of chemical bonding helps the octet rule (octet = eight). us to understandchemical formulas, e.g. why the formula for water is H20 and not H30. We will learn Octet Rule: When bonding occurs, atoms tend to how to write chemical formulas and how to name reach an electron arrangement with eight electrons compounds. in the outermost energy level. It will become quite easy for us to understandchemical bonding if we first study an important rule called the Therefore, when elements react together to form Octet Rule. compounds,their atoms tend to change their electron arrangementto try to end up with eight electrons in 5.2 The Octet Rule their outermostenergy level. It is important to realise that the octet rule is more a 'rule-of-thumb' rather than Chemists have observed that the noble gases (e.g. a strict chemical law. helium, neon, argon) in Group 0 (or 18) of the Periodic Table are very unreactive, i.e. they form practically no There are a number of exceptions to the octet rule: compounds. Hence, they are often called the 'inert' o Transition metals do not usually obey the octet rule. gases, Fig. 5.1. In many of their compounds transition metals can have more or fewer than eight electrons in their outermost energy level. CHEMISTRY LIVE  Formulas Chemical Bonding: Chemical o The elements near helium in the Periodic Table between (hydrogen, lithium and beryllium) Anioniebond is the force of attraction tend to achieve Ionic bonds the electron arrangement oppositely charged ions in a compound. of helium with two of electrons in the outer energy are always formed by the complete transfer level rather than the electrons from one atom to another. eight electrons of the other noble gases. We now study how the octet rule helps us to understand the two main types of chemical bonds —ionic bonds Common table salt (sodium chloride) is one of the best and covalent bonds. known examples of an ionic substance, i.e. a substance held together by ionic bonding. 5.3 Ionic Bonding — Transfer of Electrons How to Show the Formation Atoms can obtain the electron configuration of Ionic Bonding of a noble gas by gaining or losing electrons. When an atom gains or loses electrons it becomes a charged There are two ways of showing the formationof the atom — commonly referred to as an ion. As we shall see ionic bonding in sodium chloride: either in terms of later, a group of atoms carrying a charge is also called an the Bohr-typecircle diagram, Fig. 5.2(a), or in terms ion. of the dot-and-cross diagram, Fig. 5.2(b). An i6iiiis a charged atom or group of atoms. The elements in group I of the Periodic Table tend Na atom Cl atom Na+ ion cr ion to lose one electron to form an ion with one positive charge. (b).............+[Nap + o The elements in group Il of the PeriodicTable tend to lose two electrons to form an ion with two Sodium atom Chlorine atom Sodium ion Chloride ion positive charges. 2.8,8 o In a positive ion, since the atom has lost electrons, Fig. 5.2 (a) The formation of the ionic bond in sodium chloride using there are more protons in the nucleus than there are the Bohr-type diagram. (b) The formation of the ionic bond in sodium chloride using the dot-and-cross diagrams electrons in orbit around it. o The formationof some positive ions (also called cations) may be shown simply as follows: Note the following points: In drawing electron dot-and-cross diagrams the outer electrons of one atom are representedby dots, with crosses used for the electrons of the other atom. o The elements in groups VI (or 16) and Vll (or 17) The numbers under the chemical symbols show the of the Periodic Table tend to gain electrons to form arrangement of electrons in the energy levels. negative ions (also called anions). Cl + e- Cl- O + 2e- 02- The transfer of electrons is shown by the arrow in the dot-and-cross diagrams. 2.8,7 Both methods of representingionic bond formation Note that in each case the ion formed has the are equally acceptable in an examination provided electron configuration of the nearest noble gas. It is the electron configuration is written as shown in the formation of this stable electron configuration Fig. 5.2. Bohr-typecircle diagramscan be time that is the 'driving force' behind the formation of the consuming to draw and give more information ion. Once the ions are formed, they are held together than we need, since it is only the electrons in in the compound by the force of attractionbetween the outermost orbit that are involved in bonding. the oppositely charged ions. This force of attraction Therefore, we shall use dot-and-cross diagrams as is called an ionic bond. much as possible in this textbook. 45 Leaving Certificate Chemistry Example 5.1 Question Show the formation of the ionic bond in magnesium fluoride, Mg%, by means of a suitable diagram, Answer Each magnesium atom losestwo electronsand each fluorine atom gains only one electron. Therefore, there must be two fluorine atoms for every one magnesium atom (Fig. 5.3). Mg. Magnesium atom Fluorine atom Fluorine atom Magnesium ion Fluoride ion Fluoride ion 2.8.2 2.7 2.8 Fig. 5.3 The formation of the ionic bond in magnesium fluoride, Mg%. Sodium Chloride Crystal Structure o In Fig.5.4 we see the repeatingunit in the crystal lattice.This repeating unit is called the unit cell of X-ray studies enable chemists to work out the sodium chloride and it repeats itself in all directions arrangement of ions in a crystal. Sodium chloride, for to build up the crystal. example, has a cubic structure, Fig. 5.4. e In the crystal lattice, each sodium ion is surrounded by six chloride ions and each chloride ion is surrounded by six sodium ions. In fact, ions are so small that the tiniest crystal of sodium chloride consistsof millions of ions! Sodium chloride (commonly called 'salt') is one of the most important ionic compounds in our lives. In Roman times soldiers received part of their pay in salt- the Latin word for salt is 'sal', hence the word 'salary'. Salt is an essential part of our diet. We need to take in Na+ about 0.5 g of salt every day to replace what lose in sweat, tears and urine. However, most of us take in a lot more than this as much of the food we eat contains salt. Too little salt causes cramp in our muscles. Hence, many athletes take salt tablets to replace the salt they lose in sweating, Fib'.5.5. Too much salt in your diet over a long period can cause high blood pressure - a major cause of heart disease and strokes. Fig 5.4 Th6 regesents the crystal structureof sodium chbt&, [e. ate packed together in the crystal. Note the following points about the structure of sodium chloride: Although fig. 5.4 shows only 27 ions, remember that the structure of sodium chloride consists of a regular three-dimensional arrangement of millions of scxfium ions and chloride ions. The three-dimensional arrangement of ions is called Fig 5.5 Athietes take to to k'se sneamg. Sweatts sobJtm. a crystal lattice. 46 CHEMISTRY  used as a food preservative and in the chemical try to manufacture soap, sodium metal, chlorine detergents, toothpaste and washing Example 5.2 soda. Also, it read on roads during the winter to help melt frost snow. Question 4 How to Write the Formulas Write the formula of potassium bromide. of Ionic Compounds Answer chemical formula is a way of representinga The potassium ion is K+. The bromide ion is Br-. mpound using symbols for the atoms presentand The formula K+Br- is correct since there is the Imbers to show how many atoms of each element e present. We have seen in Example 5.1 that the same number of positive and negative charges. hemical formula for magnesium fluoride is Mg%. This Answer: The formula of potassium bromide is ells us that the compound magnesium fluoride consists KBr. magnesium and fluorine. It also tells us that there twice as many fluoride ions as magnesium ions in a crystal of magnesium fluoride. For your examination you must know how to write the Example 5.3 formulas of ionic compounds of the first 36 elements. The following are the key points: Question o Ionic compounds are usually formed between the elements of groups I and Il (metals) and those of Write the formula of calcium chloride. groups VI (or 16) and Vll (or 17) (non-metals) of the Periodic Table. In general, ionic compounds are Answer usually formed when metals react with non-metals. (The 'steps of stairs' going from boron to astatine The calcium ion is Ca 2+. The chloride ion is Cl-. divide the metals on the left from the non-metals The formula Ca 2+Cl- is not correct since there on the right.) are two positive charges but only one negative charge. To get two negativecharges we must o Metals have a tendency to lose electrons and non- metals have a tendency to gain electrons. have two Cl- ions. Answer: The formula of calcium chloride is CaClz. o An ionic compoundis neutraloverall.Therefore, there must be the same number of positive charges compound. as negative charges present in the (excluding the Some ions of the first 36 elements Example 5.4 d-block metals) are shown in Fig. 5.6. Question 111tv v vl Write the fomula of sodium sulfide. l Na Mg Art Crx Answer -V cab Ch ge blo 10 c not p odl od S2-. The sodium ion is Na+.The sulfide ion is since there is 36 elements. The formula Na+S2- is not correct Fig. 5.6 Some ions of the first charges. one positive charge but two negative must have two To get two positive charges we will help you to Studying the following examplesformulas of ionic Na+ ions. the sulfide is NazS. understand how to write Answer: The formula of sodium compounds.  Leaving Certificate Chemistry Example5S Example 5.6 Question Question Write the formula of potassium hydroxide Write the formula of aluminium oxide. Answer Answer The potassium ion is K+. The hydroxide ion The aluminium ion is AP. The oxide ion is is OH-. The formula K+OH- is correct since 0 2-. The formula A13 02- is not correct since there is the same number of positive and there are three positive charges but only two negative charges. negative charges. Bring all charges up to Answer: The formula of potassium their lowest common denominator,i.e. 6. hydroxide is KOH. To get six positive charges we need two aluminium ions. To get six negative charges we need three oxide ions, i.e. (0 2-)3 Example 5.7 Answer: The formula of aluminium oxide is A1203. Question Writing Formulas of Compounds Write the formula of sodium sulfate. with Group Ions Answer There are many ionic compoundsthat containgroup The sodium ion is Na+. The sulfate ion is ions, e.g. the sulfate ion, the carbonate ion and the nitrate ion. For example, barium sulfate is used in S042-. The formula Na+S042- is not correct X-ray examinations of internal organs, chalk is calcium since there is one positive charge but two carbonate and sodium hydrogencarbonate (bread negative charges. To get two positive soda) is used in baking. The group ions commonly charges we must have two Na+ ions. encountered in this course are listed in Table 5.1. It is Answer: The formula of sodium sulfate is not possible to predict the formulas of group ions from Na2S04. the Periodic Table —so learn them off by heart! Name Formula Sß Example Hydroxide ion OH- Nitrate ion One negative charge Hydrogencarbonate ion HC03- Question Permanganate ion Mnoa- Write the formula of calcium Carbonate ion CO 2- hydrogencarbonate. Chromate ion Cro 2- Dichromate ion C r 072— Answer Two negative charges Sulfate ion The calcium ion is Ca 2+. The Sulfite ion hydrogencarbonate ion is HC03-. The Thiosulfate ion formula Ca 2+HC03- is not correct since Ph0sphate ion PO 3- Three negative charges there are two positive charges but only Ammonium ion One positive charge one negative charge. To get two negative Table 5.1 Some common group ions. charges we must have two HCOC ions, i.e. The following examples will help you to write the Answer: The formula of calcium formulas of ionic compounds that have group ions. hydrogencarbonate is 48 CHEMISTRY LIVE IvntingFormulas of Compounds 20ntaining Transition Metals Example (ou will see in Fig. 5.6 (p. 47) that we have indicated that it is not possible to predict the charges of the Of the d-block ions elements. In this course we will Question number of meet a transition metal elements, some of which compounds.The d-block Write the formula of iron(ll) carbonate. are transitionmetals,often have variable valency, Fig. 5.7. The word valency means 'combining power'. Answer (A more formal definition of valency will be given in section 5.6.) The Roman number (Il) indicates Fe2+. The following are the main transition metals that show The carbonate ion is C03 2-. The formula variable valency and their compounds that we will Fe2+CO is correct since there are two meet in this course: positive charges and two negative charges. o Iron combines with chlorine to form either FeC12or Answer: The formula of iron(ll) carbonate is FeC13.The compound FeC12is an ionic compound and the +2 charge on the iron atom is Fecq. represented by putting the Roman numeral Il in brackets. Therefore, FeC12is called iron(ll) chloride. Similarly, the compound FeC13is called iron(lll) chloride. Example SAO Copper combines with oxygen to form either Cup or CuO. Cup is called copper(l) oxide and Cuo is Question called copper(ll)oxide. o Chromium exists as Cr3+ ions in chromium(lll) Name the compound chloride, CrC13.In Na Cro sodium dichromate, it appears that there are Cr6* ions. (We shall clarify Answer this in Chapter 14.) Hence sodium dichromateis The three sulfate ions together give a total also written sodium dichromate(Vl). negative charge of -6. Since the overall Manganese also shows variable valency. Mn02 is charge must be zero, the two chromium an ionic compound called manganese(lV)oxide — atoms must carry a total charge of +6. commonly called manganese dioxide. MnS04 is Therefore, each Cr atom must be Cr3+. called manganese(ll) sulfate and KMn04, commonly Therefore, the name of the compound is called potassium permanganate. It is also called chromium(lll) sulfate. potassium manganate(Vll). Answer: Chromium(lll) sulfate Do not get confused between compounds whose names end in -ide and those whose names end in -ate. A compound that contains just two elementsalways ends in -ide, e.g. magnesium sulfide, Mgs. (Hydroxides are an exception to this rule). Compounds whose names end in -ate contain oxygen as well as the other two elements, e.g. magnesium sulfate, MgS04. Fig. 5.7 All of these solutions contain ions of d-block elements. 5.5 d-Block Elements and All of these ions have variable form valency. coloured Note also that when solutions. Transition Elements dissolved in water they all In many cases there is no distinction made between the d-block elements and the transitionelements.When because there Transition metals exhibit variable valency referring to the s, p, d and f blocks of the Periodic Table, the 4s and is such a small energy difference between transition elements are regarded as those in the d-block lose different of the table. However, if one examines the first row of 3d sublevels. This means that they can to give metal that the numbers of electrons from these sublevels the d-block from scandium to zinc, it is found to ions with different positive charges. properties of scandium and zinc are quite different 49 Leaving Certificate Chemistry those of the other eight elements in this row. All of the In the case of Sc3+the d sublevel and in the case ofZn2+ the d elements from titanium to copper inclusive are referred subleve Since the properties of transition to as transition elements or transition metals. These me with partly filled d sublevels, show certain characteristics: scandi not considered transition metals. However In sh Transition metals have variable valency. elements from scandium to zinc forms are d scandium only forms Sc3+ ions and zinc only but only those eight elements from tita Zn 2+ ions. (inclusive)are transition metals. coloured o Transition metals usually form compounds. However, scandium and zinc only Test Yourself: Attempt qt form white compounds. 5.1 and W5.1-w 5.2 catalysts o Transition metals are widely used as (Chapter 16). However, scandium and zinc show 5.6 Covalent Bonding little catalytic activity. Sharing of Electrons Since scandium and zinc do not have typical transition definition of When studying ionic bonding we learned metal properties, chemists devised a transition metals that excludes these two elements but compounds are usuallyformed when metals re non-metals. However, there are many includes the elements from titanium to copper. compou contain only non-metals, e.g. water, ammon carbon dioxide. These compounds do not conta A transition metal is one that forms at least one ion ions. Instead, the atoms achieve noble gas configur with a partially filled d sublevel. by sharing electrons. To help you understand this of electrons being shared, we now consider a nu Neither Sc3+ nor Zn 2+ has a partially filled d sublevel. of examples. Sc3+ = 1s2, 2s2, 2p6 3s2, 3p6 A molecule is a group of atoms joined together.It is Zn2+= 1s2, 2s2, 2p 6, 3s2, 3p6 3d 10 the smallest particle of an element or compoundthat can exist independently.  The Water Molecule, H20. Lone pair Lone pair An oxygen atom has six electrons in its outer energy level. This atom needs more electrons to complete its octet. To achieve this, an oxygen atom shares two of its own electrons with two hydrogen atoms. Thus, two covalent bonds are formed as shown in Fig. 5.10. The water molecule is a V-shaped molecule. We shall understand the reason for this at a later stage. Note that there are two pairs of electrons on the oxygen atom that are not involved in bonding. These two pairs of electrons are called 'Ione pairs'. They are called 'lone pairs' because they are not bonded to another atom. Pairs of electrons that are involved in bonding (i.e. Fig. 5.10 The water molecule contains 'tied up' in covalent bonds) are called 'bond pairs'. Thus, the oxygen two covalent bonds and two lone pairs. atom in a water molecule has two lone pairs and two bond pairs.  Leaving Certificate Chemistry The valency of an element is defined as the Example 5.15 of atoms of hydrogen or any other monovalentnumb elementwith which each atom of the element combines. The Methane Molecule, CH,. A carbon atom has four electrons in its outer Therefore, since chlorine combines with one energy level. This atom needs four more hydrogen(HCI) we say it has a valency of one. atom Si electrons to complete its octet. To achieve oxygen combines with two atoms of hydrogen it h valency of two (1-120).Similarly, nitrogen has a this, a carbon atom shares four of its own vale of three(NHJ) and carbon a valency of four (CHO electrons with four hydrogen atoms. Thus, four covalent bonds are formed as shown in Fig. 5.12. Doubleand Triple Bonds In Examples 5.11—5.15 above we have seen how atoms shared single pairs of electrons to form covalent bonds x Such bonds are more correctly referred to as single covalent bonds. o A single bond is formed when one pair of electrons is shared between two atoms. o A double bond is formed when two pairs of electrons are shared between two atoms. The Fig. 5.12 The methane molecule contains four covalent oxygen molecule contains a double bond and the bonds. carbon dioxide molecule contains two double bonds, Fig. 5.14. From the above examples it is clear that one can predict the number of covalent bonds formed around certain atoms: hydrogen atoms have one covalent bond, oxygen atoms have two covalent bonds, nitrogen atoms have three covalent bonds, and carbon atoms have four covalent bonds. The number of covalentbonds formed by atoms of some selected elements is given in fig. 5.13. o Fig. 5.14 Both the oxygen molecule and the carbon dioxide molecule contain double bonds. o From fig. 5.14 we see that the formationof the double bonds enables each element to achieve the octet of electrons in the outer energy level. o A triple bond is formed when three pairs of Fig. 5.13 The numberof covalentbonds formedby atoms of some selected elements electrons are shared. The nitrogen molecule, IN? is an example of a molecule that contains a triple bond, Fig. 5.15. Since a hydrogen atom cannot combine with more than one atom of any other element, we say its valency ('combining power') is one or it is monovalent. In fact, hydrogen is used as the standard by which valency is measured.  Chemical Bonding: Chemical Formulas e The bonding in the chlorine molecule can be explained by overlap of orbitals. If we consider the Example 5.16 electron configuration of chlorine we see that one of its p orbitals is half filled. Draw a 'dot-and-cross' structure showing Cl = 1s2, 2s2, 2p6, 3s2, 3p 2, 3p 2, 3pz1 the bonding in HCHO. o The covalent bond in C12 is formed when two In showing the bonds in this molecule, half-filled pz atomic orbitals (one from each atom) rememberthat each hydrogen atom has overlap head on to form a sigma bond, Fig. 5.18. just one bond around it, the carbon atom must have four bonds around it and the oxygen atom has two bonds around it. The valencies of these atoms are satisfied by the Overlapping p-orbitals Molecular orbital structure as shown in Fig. 5.16. Fig. 5.18 The covalent bond in C12being formed by the head-on overlap of two p orbitals. There is very little difference between the orbital approach to single bond formationand the simple 'shared electron pair' approach as already discussed in Examples 5.11—5.16. Both approaches describe the same occurrence i.e. all single bonds can also be called sigma bonds. Fig. 5.16 The bonding in HCHO shown using the electron The two types of description differ when we dot method. describe the formation of double and triple bonds. Consider how the double bond in the 02 molecule is formed. Sigma and Pi Bonding O = 1S2, 2S2, 2px2, 2P l , 2P I It is also possible to describe covalent bonding in terms There are two half-filled p orbitals in the oxygen atom. of atomic orbitals. Imagine two oxygen atoms approaching each other The formation of the 1-12molecule can be described so that the two 2P orbitalsoverlap head on to form by saying that the Is orbitals of two hydrogen atoms a sigma bond, Fig. 5.19. The two pz orbitals overlap overlap as shown in Fig. 5.17. sideways to form what is called a pi (r) bond. Apjßond is formed by the sideways overlap of p orbitals. Is orbital Overlap Molecular orbital Is orbital overlap of the Fig. 5.17 The covalent bond in H2 being formed by the Is atomic orbitals of two hydrogen atoms. z z o When the two atomic orbitals overlap they form a molecular orbital. This is rather like two soap bubbles coming together to form a larger soap bubble. The covalent bond formed by the head-on overlap of the two atomic orbitalsis called a sigma bond. The Greek letter o (sigma) is used to indicate a sigma bond. Fig. 5.19 The double bond in oxygen consists of one sigma and one pi bond. A sigmmborydis formed by the head-on overlap of two orbitals. CHEMISTRY LIVE 53 Leaving Certificate Chemistry o Consider how the triple bond in the nitrogen of ionic and covalent compounds can be understoodif molecule, N2, is formed. we keep these facts in mind. N = 1s2, 2s2, 2px1, 2P 1, 2p 1 (a) Hardness There are three half-filled p orbitals in the nitrogen Ionic compounds are usually difficult to cut. This atom. Imagine two nitrogen atoms approaching each is because each ion is held in the crystal latticeby other so that the two 2pxorbitals overlap head on to strong attractive forces between it and the ions of form a sigma bond, Fig. 5.20. opposite charge around it. o Covalent compounds are usually quite soft because z they consist of molecules, e.g. iodine consistsof dark grey crystals that are so soft that they can be crushed between your fingers! A crystal of iodine consists of millions of iodine molecules, 12.Hence iodine is referred to as a molecular crystal in which z there are only very weak forces between one iodine molecule and another. (The nature of these weak forces will be studied in section 5.10.) Fig. 5.20 The triple bond in nitrogen consists of one sigma and two pi bonds. (b) Melting and Boiling Points The two 2P orbitalsoverlap sidewaysto form a pi O In ionic compounds the ions can only vibrate about fixed positions. When heat is supplied, the ions gain bond. energy and vibrate more and more. Eventually,the A second pi bond is formed from the sideways ions vibrate so much that they break away from their overlap of the two 2pzorbitals. fixed positions in the crystal. When this happens, Therefore, the triple bond in nitrogen consists of the crystal structure breaks down and we observe one sigma and two pi bonds. This is true of all triple the substance melting. However, because the forces bonds. of attraction between the ions are so strong,most In a double bond, one of the bonds is a sigma and ionic substances have high melting points and high one isa pi bond, Fig. 5.21. boiling points. For example, the melting pointof Sigma bond Sigma bond sodium chloride is 801 oc and its boiling pointis 1413 oc. At room temperature almost all ionic NrN compounds are solid. O In covalent compounds there are only weak forces Pi bond Two pi bonds between the molecules. Therefore, the melting points Fig. 5.21 A double bond consists of one sigma and one pi bond. and boiling points are much lower, e.g. the melting A triple bond consists of one sigma and two pi bonds. point of iodine is only 114 0C and its boiling pointis only 184 oc. Other covalent compounds have even o Sigma bonds are stronger than pi bonds as there lower melting and boiling points e.g. water (m.p.= is more overlapping of orbitals in sigma bonds. OOC, b.p. = 100 0C), ammonia (m.p. =-77 0c, b.p.= Therefore,an 0=0 double bond is not twice as —33oc), methane(m.p. = -182 0C, b.p. = -164 00. strong as an 0—0 single bond since pi bonds are not Therefore, at room temperature most covalent as strong as sigma bonds. compounds are usually either liquids or gases. The orbital theory gives us a detailed picture of the formation of double and triple bonds. This is the only use (c) Conduction of Electricity we will make of this theory in our study of chemistry for the Leaving Certificate. For all other cases of bond description o Ionic compounds do not conduct electricityin the the simpler 'shared-pair theory' is perfectly adequate. solid state. The reason for this is because the ions are locked in fixed positions and cannot carrythe 5.7 Characteristics of Ionic electric current through the crystal. However, if the ionic substanceis melted or dissolved in water, and Covalent Compounds the ions are now free to move. Therefore,molten We have seen that ionic compounds usually consist of a ionic compounds and ionic compounds in solution network of ions whereas covalent compounds generally conduct an electric current. The apparatusto test consist of individual molecules. Many of the properties for the presence of ions in solution is shownin Fig. 5.22. CHEMISTRYY  Chemical Bonding: Chemical Formulas covalent compounds consist of separate molecules. The molecules of covalent compounds have particular shapes. In 1940, Sidgwick and Powell proposed a theory to account for the shapes of molecules. Their theory is known as the Valence Shell Electron Pair Repulsion Theory or VSEPR Theory for short. (The term 'valence shell' simply means 'outer energy level'.) This theory states that the shape of a molecule depends on the number of pairs of electronsaround the central atom. Since electrons are negatively charged, the electron pairs repel each other and arrange themselves in space so that they are as far apart as possible. presere ot We will now consider some examples to help us understand how the VSEPR theory works. Fig. 5.22 Apparatus to test if a solution will conduct electricity.The bulb lights if ions are present in the liquid being tested. Example 5.17 o Covalent compounds do not usually conduct electricity as solids, liquids or gases or when dissolved in water. The reason for this is because covalent Beryllium chloride, BeC12 substances do not contain ions to carry the electric current. Therefore, if the beaker in Fig. 5.22 contains substances such as distilled water, alcohol, petrol or 1800 sugar solution, the bulb does not light. The propertiesof ionic and covalent compounds are summarised in Table 5.2. Ionic Covalent Contain a network of ions in Contain individual molecules. the crystal. 2. Usually hard and brittle. Usually soft. Fig. 5.23 The beryllium chloride molecule has two pairs of Have high melting points Have low melting and electrons around the central atom. The bond angle 3. and boiling points. boiling points. is 1800. Usually solid at room Usually liquids, gases or soft 4. A molecule of beryllium chloride consists temperature. solids at room temperature. of two pairs of electrons around the Conduct electricity in molten central atom, (Fig. 5.23). (The lone pairs of 5. state or when dissolved in Do not conduct electricity. electrons on the chlorine atoms themselves water. are ignored by the VSEPR theory —we need Table 5.2 The properties of ionic and covalent substances. only consider the electrons that are actually around the central atom.) By experiment it Note: The fact that the bulb lights indicates that there is found that BeC12is a linear molecule, i.e. are ions in solution. It does not necessarily mean that the bond angle is 1800. This bond angle is water the substance that was originally dissolved in the explained by saying that the two pairs of is ionic. For example, hydrogen chloride is a covalent electrons, since they are negatively charged, substance that dissolves in water to give ions: repel each other and arrange themselves HCI + Cl- 1-1+ in space as far apart as possible from each other. When the bond angle is 1800 the 5.8 Shapes of Covalent repulsion between the pairs is minimised. This shape is the shape adopted by any Molecules molecule that has two pairs of electrons In the previous section we learned that ionic around the central atom. compounds consist of giant crystal lattices whereas CHEMISTRY LIVE 55  Example 5.19. Methane CH4 A molecule of methane consists of four pairs of electrons around the central atom (Fig. 5.25(a)). When the bond angle is measured 0 experimentally, it is found that it is 109.5 , i.e. the four pairs of electronstake up the shape of a tetrahedron. By adopting this shape, the i H (c) repulsion between the four pairs of electrons is minimised. It is difficult to representthe shape of a tetrahedron on paper. The wedge drawn in fig. 5.25(b) indicates that this particular C—Hbond is coming out of the plane of the paper, the C—Hdashed line indicates the bond is going behind the paper, and the final Fig. 5.25 (a) The methane molecule has four pairs of electrons two C—H bonds are in the plane 0 of the paper. around the central atom. The bond angle is 109.5. Examining a model of a tetrahedron (b) Representing a tetrahedron on paper. (c) Model of (Fig 5.25(c)) in your school laboratory will molecule of methane. help you understandthis.  Example5.20a Question Use the VSEPR theory to deduce the shape of the ammonia molecule, NH3. Answer Step 1 Locate nitrogen in the Periodic Table. It is in group V, i.e. it has five electrons in its outer shell. Step 2 Work out the number of electrons contributed by the three hydrogen atoms to the bonds around the central atom. Since each hydrogen atom contributes one electron,the three hydrogen atoms will contribute a total number of three electrons (Fig. 5.26(a)). Fig. 5.26(a) Ammonia has the shape of a pyramid. Step 3 Work out the total number of electron pairs around the central atom = 5 electrons 3H=3x1 = 3 electrons Total = 8 electrons = four pairs of electrons tetrahedral shape. We would therefore predict that ammonia has a Fig. 5.26(b) Model of ammonia molecule. basic tetrahedral shape. By experiment, it is found that the bond angle is 1070. It is less than the regular tetrahedral bond angle of 109.50 due to the presence of the lone pair of electrons. (The ammonia molecule has one lone pair and three bond pairs of electrons.) The increased repulsion caused by the 0 presence of the lone pair causes the bond angle to decrease by about 2.50 to 107 , Fig. 5.26(a). Thus, the shape of the ammonia molecule is pyramidal (Fig. 5.26(b)). Example5.21u Question Use the VSEPR theory to deduce the shape of the water molecule. Answer Proceeding as in Example 5.20. o = 6 electrons x x = 2 electrons 17-5' Total = 8 electrons Fig.5.27 Thewatermolecule isV-shaped. = four pairs of electrons tetrahedral shape  king Certificate Chemistry We would predict that the the bond water molecule has a basic tetrahedral shape. By experiment, it is found that angle is 104.50. because of the presence of the It is less than the regulartetrahedral bond angle of 109.50 two lone pairs of two bond pairs of electrons.) electrons. (The water molecule has two lone pairs and The increased bond to decrease by repulsion caused by the presence of the two lone pairs causes the about 50 from water molecule is 109.50 to 104.50 (Fig.5.27). Thus, the shape of the V-shaped, Fig.5.27. 5.9 Electronegativity — Example 5.22 'Tug-Of-War' For Electron In a covalent bond between identical atoms (e.g. Question h-12and CID,the pair of electrons is shared equa between the two atoms in the molecule. Howevc Use the VSEPR theory to deduce the shape chemists have found that in bonds between differe, of the SiC14 molecule. atoms, the pair of electrons is often more attracte to one of the atoms than to the other. For example Answer in hydrogenchloride, HCI, it is found that the two electrons in the bond are more attracted to the Si = 4 electrons chlorine than to the hydrogen. Since the two electrons = 4 electrons spend more of their time nearer to the chlorine atom Total = 8 electrons this gives the chlorine atom a slight negative charge = four pairs of electrons and leaves the hydrogen atom with a slight positive tetrahedralshape charge. This small amount of charge is indicated by the Greek letter ö (delta), Fig. 5.28. Thereare no lone pairs present, so SiC14will have a regular tetrahedral shape. Partial Partial positive negative charge charge Example 5.23 Question Fig. 5.28 In the single covalentbond in he VSEPR HCI, the two electrons are theory to deduce the shape more attracted to the Cl atom than to the H atom. This PBGmolecule. gives the chlorine atom a slight negative charge and the H atom a slight positive charge. inswer 5 electrons The ability of an atom in a 3 electrons covalent bond to attract electrons in the bond to itself the is given a special 8 electrons and is called the name electronegativity of the atom. our pairs of electrons tetrahedralshape Electronegativityis the lone pair of electrons and three an atom in a molecule relative attraction that has for the shared electrons. The lone pair will electrons in a covalent pair of bond. i angle to about 1070. hape is pyramidal. some important points about electronegativity may be urself:Attempt o Since the questions 'electron pulling w5.3-w5.4. greaterthan that power' of chlorine is of hydrogenin chlorine is more HCI, we say that electronegati  Chemical Bonding: Chemical Formulas American chemist called Linus o A very famous A polar covalent bond is a bond in which there is Pauling, Fig. 5.29, studied the amounts of energy unequal sharing of the pair (or pairs) of electrons. needed to break certain bonds. From his results he This causes one end of the bond to be slightly set up a scale of relative values of electronegativity. positive (ö+) and the other end slightly negative (ö-). The term pure covalent is often used to refer to a covalent bond where there is equal sharing of the two electrons in the bond, e.g. there is equal sharing of the two electrons in the H—Hbond in 1-12 or the CI—CIbond in Cly Similarly, all the P—H bonds in PH3 are pure covalent since the electronegativityvalues of P and H are virtually identical (2.19 and 2.20 respectively). Uses of Electronegativity Values There are two main uses of electronegativityvalues: Fig. 5.29 In 1939 the American chemist, Linus Pauling, made an (i) To predict the polarity of covalent bonds. enormous contribution to our understanding of chemical bonding in his famous book 'The Nature of the Chemical (ii) To predict which compounds are ionic and which Bond'. In 1954 he was awarded the Nobel Prize in chemistry are covalent. and he was awarded the Nobel Prize for peace in 1963. We now explain both of these uses. For example, he found that fluorine had four 1. To Predict Polarity times the 'electron pulling power' of calcium. Therefore, he assigned an electronegativity value of Covalent Bonds of 4.0 to fluorine and 1.0 to calcium. A table of By referring to the table of electronegativityvalues, electronegativity values is shown in Fig. 5.30. we can tell if the bonds in a molecule are polar. The greater the electronegativity difference, the more polar 1 11 the bond, Fig. 5.31. 0-98 HCI CH 141 180 2-19268 316 2-20 3-44 2-20 3-16 2-55 2-20 2-20 2-20 E.N. difference = 1-24 0-96 0-35 Highly polar Polar Negligible Pure covalent covalent polar covalent covalent Fig. 5.31 The electronegativitydifferencebetween the elements in a covalent bond tells us if the bond is polar covalent or pure elements. Fig. 5.30 Electronegativity values of the covalent. electronegative o Note from Fig. 5.30 that the more as elements are the very reactive non-metalssuch In most molecules that have polar covalent bonds, the fluorine and chlorine. The least electronegative molecules themselves are also polar, i.e. the overall as elements are the very reactive metals such molecules have a partial positive and a partial negative potassium and sodium. pole which are separated by a distance, Fig. 5.32. said Elements with low electronegativity values are to be electropositive. 83- than o Since chlorine is more electronegative 6+ hydrogen, it is more correct to refer to the bond H in H—CIas a polar covalent bond. Thus we say that H hydrogen chloride is a polar covalent molecule or simply a polar molecule. The word polar simply Fig. 5.32 Hydrogen fluoride, water and ammonia have polar covälent bonds Overall, each of these molecules is a polar molecule means (a) that the bonding electrons are not shared since the positive and negative poles are separated by a equally, (b) that there is a partial positive and a partial distance. Note that since each molecule is neutral overall, negative charge on the molecule and (c) that these the sum of all the partial charges must add up to zero. partial charges are separated by some distance. 59  Leaving Certificate Chemistry in Fig. 5.33 are polar None of the molecules However, there are some molecules which, even of partial positive charge and cent since the centres r though they have polar covalent bonds, are not polar charge coincide, i.e. the of partial negative molecules. These are usually symmetrical molecules. other out due to the the bonds cancel each Studying Fig. 5.33 will help you to understandthis. of the molecules. a polar molecule gave The fact that water is chemists indication of the shape of the molecule, If an V-shaped, then it would were linear rather than o=c=o a polar molecule, Fig. 5.34. molecule were linear it would not be a polar Fig. 5.34 If the water molecule,The fact that water is V-shaped means that the partial negative charges coincides partial positive and negative charges do not cancel out and Fig. 5.33 Since the 'centre' of the charges, none of these hence water is a polar molecule. with that of the partial positive molecules are polar molecules. is Polar or Non-Polar To DemonstrateWhether a Liquid whether a liquid is A simple way of demonstrating of a negatively polar or non-polar is to study the effect a burette, charged plastic rod on a stream of liquid from Fig. 5.35 A polar liquid (e.g. water) is attracted to the charged rod. The type of charge on the rod does not make any difference to the result of the experiment. A non-polar liquid (e.g. cyclohexane) is not affected by the charged rod. If the liquid is a polar liquid (e.g. water),the streamof as the water molecules will turn so that the negative liquid is attracted to the rod shown, i.e. the positive poles of the molecules are attracted to the positively poles of the water molecules are attracted toward charged rod. the negatively charged rod. It does not matter if the negatively charged rod is replaced with a positively If the experiment is repeated with a non-polar liquid charged one. The water will still be attracted to the rod (e.g. cyclohexane), no attraction is observed since there are no polar molecules in the liquid.  Chemical Bonding: Chemical Formulas Dissolving of Ionic Compounds For example, consider the electronegativity differences of the elementsin potassiumfluoride and methane, in Water Fig. 5.37. properties of water is the One of the most important solvent. This property depends fact that it is an excellent is a polar molecule. It is found on the fact that water 0-82 3-98 2-55 2-20 and most polar covalent that most ionic substances substancesdissolve in water. An ionic substance such E.N. difference = 3-16 0-35 i.e. < 1.7 as sodium chloride dissolves in water because the ionic i.e. > 1-7 COVALENT BONDING bonding in NaCl is overcome by the strong attraction IONIC BONDING between the ions and the polar water molecules, Fig. 5.37 Electronegativity differences may be used to predict ionic fig. 5.36. These ions are dragged away from the crystal and covalent bonding. lattice and become surrounded by water molecules i.e. the salt dissolves. It is importantto rememberthat the above is not 82- a definite law but simply a rule-of-thumb which is useful for predicting the type of bonding that might 62- be expected in a given compound.There are some exceptions to the above rule-of-thumb. The main ones are lithium hydride, LiH, sodium hydride, NaH, c- potassiumhydride, KH and calcium hydride, CaH2. All 82- of these compounds are ionic and contain the hydride ion, H-. (The above rule-of-thumb would indicate cc incorrectly that the bonding is covalent). a IntramolecularBonding and Bonding Intermolecular So far we have studied the two main types of bonding— ionic bonding and covalent bonding. Since these types of bonds hold atoms together inside or within the molecule, this type of bonding is often referred to as intramolecular bonding (Latin: intra = inside Fig. 5.36 Sodium chloride dissolves readily in water. Note how the polar water molecules line up around the ions and draw or within). Think of the word INTRAvenous injection, them into solution. i.e. where the drug is injected INSIDE the vein, Fig. 5.38. We will study the dissolving of covalent compounds in water in section 5.12. 2. To Predict which Compounds are Ionic and which are Covalent in There is a useful 'rule-of-thumb' that assists us predictingwhether a compound is ionic or covalent and also the degree of polarity in a covalent bond. than O JI.7 An electronegativity difference greater indicatesioniqbonding in a compound. An electronegativity difference 'léss;than orqequabto 1.7 indicates covalent 'bonding in a compound. An electronegativity difference greater than 0.4 Fig. 5.38 An intravenous injection is so called because the drug is and less than 1.7 indicates that the covalent bond is being injected INSIDE the vein. Intramolecular bonding is bonding that occurs inside the molecule, i.e. it holds the polar covalent. molecule together. An electronegativity difference less than or equal to 0.4 indicates that the covalent bond 61  Leaving certificate Chemistry Within a bonding is bonding This temporary dipole could induce a similar that takes place dipole Covalent molecule, i.e. it holds the atoms together. in a nearby molecule. There is then an attraction bonding and polar between the opposite charges. This attraction examples of intramolecularcovalent bondingare bonding. called a van der Waals force. is There are 0ther types of bonds that exist between one VamderWaals foices are weak attractive forces molecule and between molecules resulting from the formation another. These types of bonding are referred to as intermolecular of temporary dipoles. They are the only forces of forces (Latin: inter = attraction between non-polar molecules. between). Think of the word INTERnet which allows communication BETWEEN people. Since the electronsare moving about at high speed llnteim61eculårSforces are the forces of attraction the attractionbetweenany two molecules only exists that exist between molecules. Van der Waals forces, for an instant. These van der Waals forces are very dipole—dipole forces and hydrogen bonding are weak. It has been estimated that this type Of force examples of intermolecular forces. is only about one thousandthas Strong as the covalent bond holding the atoms together in the molecule. Intramolecular forces in covalent bonds are much e van der Waals forces are the only forces of attraction Stronger than intermolecular forces between one that exist between non-polar molecules, e.g. H moleculeand another. q, N2 and Cir We now discuss three types of forces between Good evidence for the existence of these forces molecules: van der Waals forces, dipole—dipoleforces comes from the fact that gases that consist of and hydrogen bonding. non-polar molecules can be liquefied. Therefore there must be some attractive forces between the 5.10 Van Der Waals Forces molecules of non-polar substances. o It has been found that the strength of van der Waals A Dutch scientist called Johannes van der Waals (1837— 1923) carried out research into the forces between forces increases as the molecules get bigger. molecules. His work may be summarised as follows: This is because the increase in the number of electrons means bger electron clouds, allowing o Small attractive forces called van der Waals forces the temporary dipoles to form more easily. exist betweenmoleculesas a resultof temporary internal shifts in the distribution of electrons within o For example, at room temperature C12is a gas, Br2 is a molecule. a liquid and 12is a solid. This progression in boiling point is due to the increasingstrength of the van o Imagine two electronsmoving inside a molecule der Waals forces as the relative molecular masses of hydrogen, Fig. 5.39. As the two electronsmove increase, the numbers of electrons increase and inside the molecule, it may happen that, at any one the size of the electron clouds in the molecules instant, both electrons may be closer to one end increases. of the molecule than the other, i.e. one side of the molecule may be slightly negative leaving the other o Similarly, 1-12has a boiling point of —253oc but O half of the molecule slightlypositive. does not boil until —183oc. The higher boiling point is due to the fact that the van der Waals forces in oxygen are stronger than in hydrogen since the Van der Waals oxygen atom is bigger than the hydrogen atom, force i.e. it contains more electrons than hydrogen and more temporary dipoles are formed. Therefore, more heat is required to break the stronger van der Waals forces formed between the molecules Fig. 5.39 van der Waals forces are weak attractive forces between of oxygen. temporary dipoles. o van der Waals forces also exist in the case of the noble gases —all of which consist of atoms. The A temporarydipole is set up in the molecule, i.e. two boiling points of these gases increase with the size poles with one ö+ and the other ö-. It is temporary of the atoms due to the increasing strength of the because it only exists for a short time —in contrast van der Waals forces, Fig. 5.40. As you go down to the permanentdipole in a molecule of HCI. the group, the atoms increase in size and the greater  Chemical Bonding: Chemical Formulas number of electrons in the larger electron clouds allow the temporarydipoles to form more easily. (The boiling points are given on the kelvin scale. We Example 5.24 will learn about this scale in Chapter 10). Question 220 The boiling points of hydrogen and oxygen are 20.0 K and 90.2 K respectively.Account 180 for the higher boiling point of