Chapter 1: Introduction to Organic Compounds and Covalent Bonding

Learning Objectives

After completing this chapter, you are expected to be able to:

• Define organic compounds.

• Understand the difference between atoms, elements, and compounds.

• Explain how valence electrons are identified for a given atom.

• Be able to give examples of homonuclear and heteronuclear covalent bonding.

• Explain the difference between sp³, sp², and sp hybridisations in carbon chemistry.

• Explain how single, double, and triple covalent bonds in a molecule are formed.

One way of classifying chemistry as a subject is based on the nature of the compounds to be studied: ‘organic compounds’ for organic chemistry and ‘inorganic compounds’ for inorganic chemistry. We might then first ask, what makes a chemical compound organic? In life science subjects, we study the processes of life where macromolecules as chemical compounds include DNA or RNA carrying genetic information; structural proteins in muscles and the skeletal system or functional proteins as enzymes and receptors; fats and lipids including membrane structures; or carbohydrates as energy stores in the form of starch and glycogen. These macromolecules also have monomers such as amino acids for proteins, nucleotides for nucleic acids, fatty acids for fats, and simple sugars for polysaccharides (carbohydrate macromolecules). There are also thousands of chemicals involved in metabolic reactions for building (anabolism) macromolecules or for breaking them down (catabolism) such as those in the energy generation process. Such life processes are linked to what we call organic compounds, some of which are listed in Table 1.1.

One common feature of all organic compounds is that they are made from carbon. A carbon atom also bonds with other carbon atoms (the C–C bond) and/or in many cases a carbon atom is also bonded to a hydrogen atom (the C–H bond). The oxides and sulfides of carbon (carbon dioxide, carbon monoxide, carbonic acid, and carbonyl sulfide) without these C–C or C–H bonds are considered inorganic (Table 1.1). Another important feature of organic compounds is the C–C and/or C–H bonding that we call covalent bonding. Many biological molecules such as proteins, sugars, and nucleic acids are soluble in water, but other organic compounds are less soluble in water and rather dissolve in organic solvents.

Many mineral salts such as metallic chlorides, sulfates, hydroxides, nitrates, phosphates, etc. are highly polar water-soluble chemical compounds. These compounds are considered inorganic – they lack carbon and mostly involve ionic bonding. The list of inorganic compounds (see Table 1.1) is large, however, and includes many water-insoluble chemical compounds, and covalent bonding may also be involved in bringing elements together to form chemical compounds. Hence the main difference between organic and inorganic compounds is the C–C and C–H covalent bonding that we are addressing in this chapter. Before we go into the details of covalent bonding, however, we need to revise the distinction between elements and molecules, although our focus will remain on the chemistry of carbon.

In biological molecules, many inorganic ions may also be incorporated. For example, we have proteins that we call metalloproteins because they contain one or more metal ions tightly bound to their amino acid side chains. Many enzymes also incorporate metal ions while other proteins such as haemoglobin and myoglobin incorporate iron ions. Chlorophyll is an example of an organic compound that incorporates a magnesium ion.

An element is a substance that contains one type of atom. The Periodic Table of the elements shown at the beginning of this chapter lists over 100 elements. These elements are the building blocks for all substances that we encounter. An atom is the smallest amount or particle of an element. The physical and chemical behaviour of an element at the smallest unit is therefore represented by its atoms.

Two or more atoms join together chemically to form a chemical compound. The smallest unit of this chemical combination product is a molecule. This bonding could be between two atoms of the same element to form diatomic molecules such as hydrogen (H₂), oxygen (O₂), nitrogen (N₂), chlorine (Cl₂), bromine (Br₂), or iodine (I₂). Alternatively, a molecule is made when two or more different atoms are involved in chemical bonding, such as hydrogen and oxygen to form water (H₂O) or carbon and hydrogen to form methane (CH₄). In organic compounds, the molecules may be made from combinations of more than two elements. As atoms are the smallest units of elements, molecules represent the smallest units in chemical compounds. The physical and chemical properties of a given organic compound are therefore represented at the smallest possible unit as a molecule.

The Periodic Table shows the elements grouped vertically in columns called groups and horizontally in rows called periods. These groupings provide valuable information about the elements in terms of their reactivity. The Periodic Table also shows metals, non-metals, and semimetals. It further gives information on atomic and subatomic structures. All atoms are composed of a nucleus, which contains protons and neutrons, and electrons circulating around the nucleus. The mass number of an atom is the sum of the protons and neutrons, which is represented by a superscript number, along with the proton number (subscript) for each atom. This atomic notation for carbon, $612$C , shows that its proton number is 6 and the sum of neutrons and protons or mass number is 12. Thus, carbon also has six neutrons. For oxygen, with a mass number of 16 and a proton number of 8, its atomic notation is represented as $816$O, also indicating the presence of eight neutrons in its nucleus.

The proton number is also called the atomic number and is a fixed number for a particular element. On the other hand, atoms of an element may have different numbers of neutrons and are called isotopes. Some elements such as hydrogen exist mainly as one isotope; nearly all (natural abundance ∼99.98%) hydrogen atoms exist with a mass number of 1 ($11$H) but traces of mass number 2 ($12$H) and 3 ($13$H) also exist. For carbon atoms, there are far more isotopes but the predominant one is $612$C and others such as $613$C and $614$C are useful in experimental studies. Because of the presence of isotopes, the mass number of an element is presented as the average calculated based on their natural abundance. For example, for carbon, although usually presented as $612$C, the actual mass may be given as 12.0101 and hydrogen as 1.00794. Some isotopes are called radioisotopes because their nuclei are unstable and they release excess energy through spontaneous radiation such as alpha, beta, and gamma rays. In doing so, the element is transformed into a more stable form. For carbon atoms, $614$C is a radioisotope that occurs naturally and remains for thousands of years without breaking down. Its half-life is said to be ∼5730 years and hence it has uses in geological survey studies. Radioisotopes can also be produced in reactors and their main use is to generate energy, for example in the form of electricity. Radioisotopes are also used extensively in medicine for both diagnosis and therapy.

We adopt here a simple model of electronic configuration that allows us to understand easily the bonding interactions between atoms, as quantum mechanics is beyond the scope of this book. Molecular bonding, be it formed through ionic or covalent interactions, is a function of electrons and the useful information about the number of electrons for a given element comes from the proton number shown in the Periodic Table. Protons in the nucleus are positively charged and the negatively charged electrons circle around the nucleus. Looking into the three-dimensional structure of the nucleus, we should assume that the electrons occupy different shells or energy levels, with the energy level (n) dictating the maximum number of electrons possible in the shell as 2n². The first energy level therefore has a capacity of 2, the second shell 8, the third shell 18, and the fourth shell 32 electrons. At this stage, we assume that there are large energy differences between these levels – the energy increases as we move from the first to the second level and so on. We will learn later that there are also subshells or orbitals that electrons occupy within these energy levels.

Let us consider group 14 elements of the Periodic Table: carbon, silicon, and germanium. Their atomic numbers (the same as the electron number) and their electron configurations are given in Table 1.2. For carbon, with a total of 6 electrons, the first energy level (shell) is occupied with the maximum possible 2 electrons, and the remainder go into the second energy level, which has a capacity of 8. For silicon, with 14 electrons, both the first and second shells are filled with the maximum possible number of electrons allowed (2 and 8, respectively), and the remaining 4 go into the third shell. Germanium, with 32 electrons, has the third shell filled with the maximum possible 18 electrons and the remaining 4 are in the fourth shell. Since all these elements have four electrons in their outermost shells, they were previously called group IV (now group 14) elements. On this basis, we can now see that group 1 elements have one electron in their outer shell, group 2 elements have two, and so on.

You may notice from Table 1.2 that the electronic arrangement of potassium is 2.8.8.1, not 2.8.9. This is defined by what we call the octet rule, referring to the tendency of atoms to prefer to have eight electrons in the outermost shell. Hence, if the third shell is not filled with 18 electrons, the tendency is to fill it with eight electrons and move the remaining electrons to the next shell. Another good example is calcium (Table 1.2).

The electrons in the outermost shell, the number of which governs the group name, are called valence electrons. Hence the number of valence electrons of a particular element is its group number in the Periodic Table. The reactivity of the element towards other elements is dependent on these valence electrons. As mentioned above, the octet rule states that the preferred number of electrons in the outer shell is eight. Group 18 elements already meet this requirement, and the elements are inert or unreactive. Hence neon (Ne) and argon (Ar) are also known as inert gases. The same applies to helium (He), where the two electrons fully occupying the first shell make it unreactive. On this basis, we can now see the tendency of elements to either lose or gain electrons to fulfil this rule. Group 17 elements with seven valence electrons need only one more electron to fill them up and prefer to gain one rather than lose seven. On the other hand, group 1 elements with one valence electron tend to lose this electron. Hence group 1 elements have good metallic characteristics with a tendency to lose an electron, and group 17 elements (also called the halogen series) have good non-metallic characteristics with a tendency to gain an electron. How this works out in the formation of chemical bonds is discussed in the following sections. We can represent valence electrons as dots around the atomic symbol (see the following sections and Figure 1.1).

Having the valence electrons and the various groups of the Periodic Table explained, we can also make some generalisations about the periods – the horizontal rows in the Periodic Table. You might ask what is common for period 1 elements (H and He), what is common for all period 2 elements (Li, Be, B, C, N, O, F, and Ne), and so on. If we consider group 1 elements (first column of the Periodic Table), moving down from hydrogen (H) to francium (Fr), the atomic size increases. In electronic configuration, period 1 elements have one shell whereas period 2 elements have two shells. As the numbers of electrons and shells increase, the tendency for electrons to be lost (instead of gained) or the metallic character increases. Note that a larger atomic size (greater atomic radius) means an increased distance from the nucleus (positively charged protons) that hold electrons, i.e. a greater tendency to lose an electron or an increased metallic character. Hence the metallic character increases as we move from right to left horizontally and from top to bottom vertically in the Periodic Table.

Before we move to covalent bonding, let us investigate ionic bonding, which occurs through the interaction between metal (e.g. sodium, Na) and non-metal (e.g. chlorine, Cl) atoms to make a salt (e.g. sodium chloride, NaCl). Sodium with 11 electrons has a 2.8.1 electronic arrangement whereas chlorine with 17 electrons has a 2.8.7 electronic arrangement. Since the third shell has a capacity of 8 electrons, sodium with 1 electron (as a group 1 element of the Periodic Table) favours losing the single electron in the outer shell rather than trying to gain 7 electrons, whereas chlorine as a group 17 element favours gaining 1 electron instead of losing 7 (Figure 1.1).

The electron(s) in the outer shell, which is the furthest from the positively charged nucleus, may be either lost or gained to establish stability. This results in Na losing an electron and Cl gaining an electron as follows:

Na – e^– → Na⁺ (sodium loses an electron and is positively charged)

Cl + e^– → Cl⁻ (chlorine gains an electron and is negatively charged)

The resulting ions and attractions of oppositely charged particles to make ionic bonds is shown in Figure 1.2.

The strong attraction between these two oppositely charged particles forms the ionic bond. Chemical compounds or molecules formed through ionic interactions make a uniform solid lattice or crystal structure with a high melting point. The solid lattice structure, however, crumbles when water is added to salts as the water molecules surround the two ionic species such as Na⁺ and Cl⁻, i.e. salts dissolve in water. In later sections of this book, we will address how to make organic compounds water soluble by changing them into a salt form – meaning changing them into charged species and making ionic bonds. This has implications for how drug molecules are administered orally in aqueous solutions, how they are made, how they are absorbed from the gut, and how they interact with biological targets.

In ionic bonding, we have seen how atoms gain stability through bonding by either losing or gaining an electron or electrons from their outer shells. Ionic bonding works perfectly with a combination of atoms highly deficient in electrons in their outer shells (metals) and an electron excess (non-metals), particularly group 17 elements. In other elements, this ‘give and take’ may not be appropriate to achieve stability and they rather form bonds by sharing some electrons. This is what we call covalent bonding, and it is a common feature of bonding between non-metallic elements.

As in ionic bonding, the electrons involved in covalent bonding are in the outer shells of the atoms. The goal is still to gain a stable form of a molecule with complete electron sets in the outer shell. Thus, in the simplest example of covalent bonding, we expect a pair of electrons shared between two atoms to make a bond. The number of bonds involved depends on the number of electrons available for the sharing process in the outer shell. Let us consider the first elements from group 14 to group 17 of the Periodic Table as shown in Table 1.3.

Carbon, with four electrons in the outer shell, which has a capacity of eight, needs four more electrons to achieve a stable atom in the formation of a molecule through bonding with other elements. This means that it must share two pairs of electrons – its valency is four and it can make four bonds. Oxygen (group 16 element) always has a valency of two as it needs two electrons to fill its shell, whereas group 17 elements, as shown in Table 1.3 for fluorine, have a valency of one.

Perhaps the simplest form of covalent bonding to visualise is that between the same atoms to form diatomic molecules. Elements that form covalent bonds through this process are shown in Table 1.4. Since two atoms of the same element are involved in the bonding process, a homonuclear molecule is the terminology used to describe a diatomic molecule. In single bond formation, one pair of electrons is shared as shown for hydrogen (H₂) and the group 17 diatomic elements fluorine (F₂), chlorine (Cl₂), bromine (Br₂), and iodine (I₂). For the group 16 element oxygen, two pairs of electrons contribute to the diatomic molecular bonding with two or double bonds. Nitrogen (N₂), as a group 15 element, needs three pairs of electrons in the sharing process to make a three- or triple-bonded molecule (see Figure 1.3).

Using the diatomic model of atoms forming a molecule through covalent bonding, let us consider hydrogen (H) and chlorine (Cl) atoms forming hydrogen chloride (HCl). As shown in Figure 1.4, hydrogen is a group 1 element with just one electron in its shell whereas chlorine, as explained in the previous section, has seven electrons in its outer shell. Both elements need one electron to complete their shell and hence make a diatomic heteronuclear molecule, hydrogen chloride.

We have already explained that the electronic arrangement for atoms is based on energy levels or shells with the maximum capacity of electrons in the shell being 2n². We also have subshells or orbitals that electrons occupy within these energy levels. These are what we call the s, p, d, and f orbitals. The maximum numbers of electrons in these subshells are as follows:

• an s orbital is occupied by 2 electrons

• a p orbital is occupied by 6 electrons

• a d orbital is occupied by 10 electrons

• an f orbital is occupied by 14 electrons.

  In a simplistic presentation, the first four energy levels or shells for the indicated maximum capacity can be presented as shown in Table 1.5.

  We can indicate the number of electrons occupying the orbitals as follows:

• 1s² means 2 electrons occupying the s orbital in the first shell.

• 2s² means 2 electrons occupying the s orbital in the second shell.

• 3p¹ means 1 electron occupying the p orbital in the third shell.

• For neon (Ne) with 8 electrons (2 in the first shell and 6 in the second shell), 1s²2s²2p⁶.

• For sodium (Na) with 11 electrons (2 in the first shell, 8 in the second shell, and 1 in the third shell), 1s²2s²2p⁶3s¹.

• For fluorine (F) with 9 electrons (2 in the first shell and 7 in the second shell), 1s²2s²2p⁵.

For covalent bonding involving carbon atoms, or organic compounds, we focus only on the s and p orbitals. The s orbital is spherical and has a maximum capacity for a pair of electrons, whereas p orbitals have a dumbbell shape with a capacity for a pair of electrons. For the three p orbitals of the second shell (2p subshells), we have x, y, and z orientations, as shown in Figure 1.5. These 2p_x, 2p_y, and 2p_z subshells hold a maximum of six electrons (3 × 2 electrons for each subshell). The s orbitals in the various electronic shells are denoted 1s, 2s, 3s, 4s, … , etc., whereas p orbitals are denoted 2p, 3p, 4p, 5p, … , etc.

Let us consider the electronic configuration of chlorine, which has 17 electrons. As already highlighted in the preceding section (see Figure 1.4), the electrons are arranged as a 2.8.7 order of shells. This can be written as 1s²2s²2p⁶3s²3p⁵. The three p orbitals in the second shell are full (six electrons) whereas only two p orbitals in the third shell are full, leaving one p orbital with an unpaired electron. This can be presented as shown in Figure 1.6.

Although not relevant to carbon, we also have other orbitals such as d and f and their maximum electron occupancy for each shell is shown in Figure 1.7. Given that the energy levels of these orbitals are different, for elements with a large number of protons (or electrons) the order of arrangement is as shown in Figure 1.7 in the direction of the arrows. The orbitals of lowest energy fill first and follow the order shown by the arrows.

In the previous example of covalent bonding using a homonuclear diatomic molecule, we considered hydrogen with one electron in the 1s orbital. A hydrogen molecule (H₂) is therefore made when two 1s orbitals overlap with each other to allow the sharing of the electrons as shown in Figure 1.8.

The single covalent bond formed by sharing of the electron pair in the overlapped s orbitals is a sigma (σ) bond. In the previous example, we also mentioned that two chlorine atoms form the diatomic chlorine molecule. As shown in Figure 1.6, the electron configuration of chlorine shows one unpaired electron in one of the 3p orbitals. This means that the sigma bonding in diatomic chlorine molecules is formed by the overlapping of two p orbitals. Sigma bonding also occurs between two different atoms, as we have seen for hydrogen and chlorine combining to form hydrogen chloride. In this case, the s orbital of the hydrogen atom overlaps the 3p orbitals of chlorine to form a sigma bond (Figure 1.9).

Carbon has six electrons with an arrangement of its electrons in the two shells as 2.4, i.e. it has a valency of four and can make bonds by sharing them with another carbon atom or other elements. Through covalent bonding, it can form single, double, or triple bonds, which can be explained better by looking into hybrid orbitals or orbital hybridisation. Let us first consider sp³ hybridisation that allows carbon to undergo four sigma (or single) bonding.

Ground-state carbon has a 1s²2s²2p² configuration whereas the excited state has a promotion of one electron from 2s to a higher energy p orbital, leading to a configuration of 1s²2s¹2p³. In sp³ hybridisation, a lower energy orbital of four sp³-hybridised orbitals can be achieved, i.e. 2s and 2p orbitals can combine to give hybrid orbitals that we call sp³ (Figure 1.10), i.e. s + ppp = sp³.

Note that in sp³ hybridisation, there is no unhybridised p orbital, i.e. we have four sp³ orbitals that take a tetrahedral geometry about the carbon nuclei. One unique feature of sp³-hybridised orbitals in carbon is the bond angle of the orbitals from each other of 109.5°, i.e. a tetrahedral geometry arrangement with a bond angle of 109.5° is a characteristic feature of sp³ hybridisation in carbon (Figure 1.11).

The head-to-head overlapping of these sp³-hybridised atomic orbitals with other orbitals from another atom forms sigma bonding. Let us consider carbon bonding with hydrogen atoms in an sp³ fashion to make methane (CH₄). Figure 1.12 shows this bonding, and more examples of this in carbon chemistry are discussed in the following chapters.

The principle behind sp² hybridisation is depicted in Figure 1.13. In this case, a 2s orbital and two 2p orbitals are involved to form sp² hybridisation. This leaves one p orbital unchanged that appears perpendicular to the plane of the three sp² orbitals (Figure 1.14). The sp² orbitals on the plane are separated from each other by 120°.

In sp² hybridisation, the unhybridised p orbitals of the two carbons in ethene form pi (π) bonds, i.e. the p atomic orbitals overlap side-to-side to share an electron pair in the space just above and below the σ bond between the two carbons. This is what we call sideways overlapping. Hence, in the sp²-hybridised carbon system, we have three sigma bonding, which is similar to sp³-hybridised orbitals and one π bond. This is exemplified by ethene (C₂H₆) in Figures 1.14 and 1.15.

In sp orbitals, a 2s orbital and only one of the 2p orbitals of carbon hybridise (Figure 1.16). This forms two hybrid orbitals, leaving two p orbitals untouched. The two p orbitals lie at a right-angle to each other whereas the sp orbitals are separated from each other by 180° (Figure 1.17).

Let us consider ethyne (acetylene) as a classic example of sp hybridisation in carbon. Through head-to-head overlap of the sp orbitals from each carbon, a σ bond is formed, whereas each carbon also makes a σ bond with a hydrogen atom (Figure 1.18). The unhybridised two sets of p orbitals form two π bonds between the carbons. As a result, a triple bond between the carbons is formed. The characteristic feature of sp-hybridised carbon is the 180° bond angle of the H–C≡C axis (Figure 1.18).

A good summary of orbital hybridisation in carbon is shown in Table 1.6. We have now seen examples of carbon bonding with hydrogen or another carbon atom through these three forms of orbital hybridisation. Carbon also makes bonds with various other elements and the rule of hybridisation still applies in its covalent bonding with non-metal atoms such as nitrogen, sulfur, and oxygen. We will address more of this topic in the following chapters.

In C–C bonding, the various combinations of hybrid orbitals are as summarised in Figure 1.19. A molecule may possess only C–C single bonds that are formed by head-to-head overlapping of sp³–sp³ orbitals to make σ bonding. A C=C double bond has a sigma bond through overlapping of sp²–sp² orbitals and a C≡C triple bond has a sigma bond component through overlapping of sp–sp orbitals. Molecules that possess a C=C (sp²–sp²) bond may also have sp²–sp³ σ bonding (Figure 1.19). The same applies for a CC triple (sp–sp) bond, which could undergo σ bonding with either an sp² or an sp³ carbon (Figure 1.19). Complex organic compounds and their systematic classifications are addressed in the following chapters.

In this book, the emphasis is to scrutinise the basic chemistry of organic compounds and drug molecules. One common feature of all organic compounds is that they contain carbon, which forms bonds with other elements through covalent bonding. In this chapter we have answered the following questions:

• What are organic compounds and in what way do they differ from inorganic compounds?

• What are the distinguishing features of atoms, elements, and compounds?

• What are isotopes?

• What are valence electrons and what is the valency of an element?

• What are shells and subshells/orbitals and what are their maximum capacities for electrons?

• What are groups and periods in the Periodic Table?

• How is an ionic bond formed?

• How is a covalent bond formed?

• What are homonuclear and heteronuclear diatomic molecules?

• What are the distinguishing features of sp³, sp², and sp hybridisations?

1. If the number of neutrons in an element is 7 and the number of electrons is 6, what is the atomic mass?

2. The atomic mass of hydrogen is 1.00794 not 1. Explain.

3. How many neutrons are in a carbon-14 (¹⁴C) isotope?

4. If the natural abundances of ¹²C and ¹³C are 98.9 and 1.1%, respectively, what is the fractional abundance of the two forms and their fractional sum?

5. Atoms have no overall electrical charge – why?

6. What is the valency of an element with the electronic configuration 2.8.6?

   Questions 7–11 are based on Figure 1.20.

7. Indicate the hybridisation state for the six carbons shown in structure 1. What would be the bond angle at each carbon?

8. What is the hybridisation of each carbon atom in benzene (structure 2)?

9. What is the hybridisation state of each carbon in structure 3? What would be the bonding geometry?

10. How many π bonds are in the molecule shown as structure 2 or benzene?

11. How many sigma bonds are present in structure 4?

12. When potassium loses an electron to become a potassium ion (K⁺), its electronic configuration becomes the same as that of which inert gas?

13. When a covalent bond is formed between two atoms, what happens to their valence electrons.

14. The number of bonds that an element can form is called what?

15. Complete Table 1.7.

16. What is the difference between a σ bond and a π bond in carbon?

1. The number of neutrons is given as 7 and the number of electrons given as 6:

   • number of electrons = number of protons = 6;

   • atomic mass = number of neutrons + number of protons = 6 + 7 = 13.

2. Hydrogen has three isotopes: protium ($11$H), deuterium ($12$H), and tritium ($13$H). The predominant isotope is protium and calculation of their natural abundances gives the average mass unit as 1.00794, i.e. the sum of the atomic masses of its naturally occurring isotopes, with each one multiplied by their respective abundance.

3. Carbon has six protons and therefore has six neutrons in ¹²C, seven in ¹³C, and eight in ¹⁴C.

4. An abundance of 98.9% for ¹²C means a fractional abundance of 0.989 × 12 = 11.868:

   • a 1.1% abundance of the ¹³C means a fractional abundance of 0.011 × 13 = 0.143;

   • 11.868 + 0.143 = 12.011. Note that this is an exemplary calculation and does not take into account the other isotopes of carbon.

5. There must be a balance between the positively charged protons and the negatively charged electrons. Hence atoms must have equal numbers of protons and electrons.

6. A valency of two needs two electrons to be gained to make up a stable eight electrons in the shell.

7. For structure 1 (Figure 1.20):

   • carbon 1 sp²; a carbon double bond, whether it is a bond with carbon or oxygen, is still sp²; bond angle 120°;

   • carbon 2 sp; bond angle 180°;

   • carbon 3 sp; bond angle 180°;

   • carbon 4 sp³; bond angle 109.5°;

   • carbon 5 sp²; bond angle 120°;

   • carbon 6 sp²; bond angle 120°.

8. All carbons in benzene are sp² hybridised.

9. All carbons in the molecule are sp hybridised. The molecule is therefore linear as all the bonds have a 180° bond angle.

10. Structure 2 (benzene): three double bonds mean three pi bonds.

11. Six sigma bonds. Note that a double bond has one sigma bond and one pi bond.

12. Potassium has 19 electrons and losing one leaves 18 electrons, which is of the same configuration as argon (Ar).

13. A covalent bond is formed by sharing some of the valence electrons. The valency of the element tells us how many electrons are involved in the sharing process.

14. Valency.

15. The answers are given in Table 1.8.

16. A sigma bond is formed by end-to-end overlapping of two hybridised s or p orbitals. A π bond is formed by sideways overlapping of p orbitals.