E-Content Presentation Subject : Chemistry Class : +1 Topic : Chemical Bonding & Molecular Structure Prepared By : Mr. Rohit Kashmiri, PGT Chemistry, J.N.V. Pandoh, Distt. Mandi, H.P. Chandigarh Region Technical Support : Mr. Rajeev Sharma Provided By F.C.S.A. J.N.V. Paprola, Distt. Kangra, H.P.
Linus Pauling won the Nobel Prize in Chemistry in 1954 for his work on chemical bonding. INTRODUCTION
A chemical bond is the physical phenomenon of chemical substances being held together by attraction of atoms to each other through sharing, as well as exchanging, of electrons or electrostatic forces. In general, strong chemical bonds are found in molecules, crystals or in solid metal and they organize the atoms in ordered structures. Weak chemical bonds are classically explained to be effects of polarity between molecules which contain strong polar bonds. Some very weak bond-like interactions also result from induced polarity London forces between the electron clouds of atoms, or molecules. Such forces allow the liquification and solidification of inert gases. At the very lowest strengths of such interactions, there is no good operational definition of what constitutes a proper definitional "bond".
Learning Objectives • Chemical Inertness of Noble Gases. • Types of Chemical Bonds. • Covalent Bonds. • VSEPR Theory & Shapes of Molecules. • Polarity in Covalent Bonds. • Quantum theory of Covalent Bonds: Valence Bond Theory. • Molecular Orbital Theory. • Hybridisation. • Hydrogen Bonding
Entry Behaviour • Why do certain atoms combine to form molecules whereas other do not? • What is the nature of the forces which hold the atoms together in molecules? • Why do atoms have fixed combining capacity?
In theory, all bonds can be explained by quantum theory, but in practice, chemical bonds are divided in several categories. Simplifications of quantum theory have been developed to describe and predict the bonds and their properties. These theories include octet theory, valence bond theory, orbital hybridization theory, VSEPR theory, ligand field theory and LCAO -method. Electrostatics and other physical theories are used to describe bond polarities and the effects they have on chemical substances. However, in combination, they constitute a powerful theory, which can be applied in almost all of chemistry.
In quantum mechanics, in simplified terms, electrons are located on an atomic orbital (AO), but in a strong chemical bond, they form a molecular orbital (MO). In many theories, these are divided in bonding, anti-bonding, and non-bonding orbitals. They are further divided according the types of atomic orbitals hybridizing to form a bond. These orbitals are results of electron-nucleus interactions that are caused by the fundamental force of electromagnetism. Chemical substances will form a bond if their orbitals become lower in energy when they interact with each other. Different chemical bonds are distinguished that differ by electron cloud shape and by energy levels.
Vocabulary Glossary of Terms • Atomicity of a gas: The number of atoms present in the molecule of a gas is called its atomicity. • Bond dipole moment ( m ).:A covalent bond between two atoms of different elements is called a polar covalent bond . A polar bond is partly covalent bond and partly ionic. The percentage of ionicity in a covalent bond is called percentage ionic character in that bond . The ionic character in a bond is expressed in terms of bond dipole moment ( m ). • BORN-HABER CYCLE: This thermochemical cycle was devised by Born and Haber in 1919. It relates the lattice energy of a crystalline substance to other thermochemical data. The Born-Haber cycle is the application of Hess's law to the enthalpy of formation of an ionic solid at 298 K.
4. Chemical bond : The chemical force which keeps the atoms in any molecule together is commonly described as a chemical bond. 5. Chemical compounds :Compounds are generally called chemical compounds because they are formed due to the chemical combination of the combining element. 6. Covalent bond : The Bond formed by Mutual sharing of electrons between the combining atoms of the same or different elements is called covalent bonds. 7. Double covalent bond :The bond formed between two atoms due to the sharing of two electron-pairs is called a double covalent bond or simply a double bond. It is denoted by two small horizontal lines (=) drawn between the two atoms, e.g., O = O, O = C = O etc. 8. Electronegative or nonmetallic character :The tendency of an element to accept electrons to form an anion is called its non metallic or electronegative character. 9. ELECTRONEGATIVITY :The relative tendency of an atom in a molecule to attract a shared pair of electrons towards itself is termed its electronegativity.
10. Electronic configuration:The distribution of electrons amongst various energy levels of a atom is termed its electronic configuration 11. HYBRIDISATION :The process of mixing of the atomic orbitals to form new hybrid orbitals is called hybridisation. 12. Hybrid orbitals :According to the concept of hybridisation, certain atomic orbitals of nearly the same energy undergo mixing to produce equal number of new orbitals. The new orbitals so obtained are called hybrid orbitals. 13. Hybridisation in carbon :Carbon shows sp3 hybridisation in alkanes, sp2 hybridisation in alkenes and sp hybridisation in alkynes. 14. HYDROGEN BOND :The bond between the hydrogen atom of one molecule and a more electronegative atom of the same or another molecule is called hydrogen bond. 15. Ionic (or Electrovalent) bond :An ionic (or electrovalent) bond is formed by a complete transfer of one or more electrons from the atom of a metal to that of a non-metal. 16. LATTICE ENERGY (Lattice Enthalpy) :The strength of binding forces in solids is described by the term lattice enthalpy ( D LH ) (earlier the term lattice energy was used). The molar enthalpy change accompanying the complete separation of the constituent particles that composed of the solid (such as ions for ionic solids and molecules for molecular solids) under standard conditions is called lattice enthalpy ( D LH° ). The lattice enthalpy is a positive quantity.
17. Lewis Formula (or Electronic Formula) of a Compound :The formula showing the mode of electron-sharing between different atoms in the molecule of a compound is called its electronic formula or Lewis formula. 18. Metallic crystals : In metallic crystals, the valence electrons of all the atoms form a pool of mobile electrons. The nuclei with their inner electrons (called Kernels) are embedded into this pool of free electrons. Thus, the constituent particles in a metallic crystal are the positive kernels in a pool of electrons. 19. NON-POLAR COVALENT BOND :When a covalent bond is formed between two atoms of the same element, the electrons are shared equally between the two atoms. In other words, the shared electron-pair will lie exactly midway between the two atoms. The resulting molecule will be electrically symmetrical, i.e ., centre of the negative charge coincides with the centre of the positive charge. This type of covalent bond is described as a non-polar covalent bond. The bonds in the molecules H 2 , O 2 , Cl 2 etc., are non-polar covalent bonds. 20. OCTET RULE : According to this theory, the atoms tend to adjust the arrangement of their electrons in such a way that they ( except H and He ) achieve eight electrons in their outermost shell. This is known as the octet rule. 21. Pi ( p ) Bond : A covalent bond formed between the two atoms due to the sideways overlap of their p -orbitals is called a pi (p) bond.
22. POLAR COVALENT BOND :When a covalent bond is formed between two atoms of different elements, the bonding pair of electrons does not lie exactly midway between the two atoms. In fact, it lies more towards the atom which has more affinity for electrons. The atom with higher affinity for electrons, thus, develops a slight negative charge, and the atom with lesser affinity for electrons a slight positive charge. Such molecules are called polar molecules.The covalent bond between two unlike atoms which differ in their affinities for electrons is said to be apolar covalent bond. 23. RESONANCE : When a molecule is represented by a number of electronic structures such that none of them can exactly describe all the properties of the molecule, but each structure has a contribution to it, then the molecule is termed as a resonance hybrid of all these structures. Such structures are called resonance structures and such a phenomenon is called resonance. 24. Resonance hybrid :When a molecule is represented by a number of electronic structures such that none of them can exactly describe all the properties of the molecule, but each structure has a contribution to it, then the molecule is termed as a resonance hybrid of all these structures.
25. Sigma ( s ) Bond : A covalent bond formed due to the overlap of orbitals of the two atoms along the line joining the two nuclei (orbital axis) is called sigma (s) bond. 26. Single Covalent Bond : A covalent bond formed by mutual sharing of one pair of electrons is called asingle covalent bond, or simply a single bond. A single covalent bond is represented by a small line (-) between the two atoms. 27. Triple covalent bond : Bond formed due to the sharing of three electron-pairs is called a triple covalent bond or simply a triple bond. 28. Valence electrons : Valence is one of the most important chemical property of the elements. The chemical behaviour of an element depends upon the number of electrons in the outermost shell of its atom. The electrons present in the outermost shell are called valence electrons. The electrons in the outermost shell are called valence electrons because the electrons in the outermost shell determine the valence of an element . 29. Valency : The combining capacity of an atom of an element is described in terms of its valency. It may be defined as, The number of hydrogen or chlorine or double the number of oxygen atoms which combine with one atom of the element is termed its valency. It may also be defined as, The number of electrons which an atom loses or gains or shares with other atoms to attain noble gas configuration is termed its valency.
Sub Topic 1: Chemical Inertness of Noble Gases Noble gases (He, Ne, Ar, Kr, Xe, Rn), do not form compounds neither among themselves, nor with other elements. Xenon however, forms fluoride and oxyfluoride compounds under drastic conditions. Now, the question is that why do these elements not show chemical reactivity? To answer this, let us consider the electronic configurations of these elements. The electronic configurations of these elements are given below.
All noble gases have their outermost shells completely filled. The atoms of all other elements which show chemical reactivity have less than eight electrons in their outermost shell. Thus, it appears that the chemical reactivity of any element is related to the numberanddistribution of electrons in its atom. Theoretical physicists have shown that certain number of electrons in certain definite energy levels give stable atoms. The atoms having a total of 2, 10, 18, 36, 54, and 86 electrons are found to be the most stable. These electronic configurations incidently correspond to those of the noble gases. So, it is because of their stable electronic configurations that the noble gases show no chemical reactivity. Logically, it means that the atoms of all other elements are not so stable, and they tend to gain stability by acquiring an electronic configuration of the nearest noble gas element. The atoms tend to adjust the arrangement of their electrons in such a way that they (except H and He) achieve eight electrons in their outermost shell. This is known as the octet rule.
VALENCE ELECTRONS AND VALENCY : The electrons present in the outermost shell (generally termed as valence shell) are called valence electrons. The number of valence electrons in the atoms of certain elements are given below. "The number of electrons which an atom loses or gains or shares with other atoms to attain noble gas configuration is termed its valency."
ELECTRONIC THEORY OF VALENCY : The combining tendency of atoms was explained by Kossel and Lewis (1916) through their theory called electronic theory of valency. The main postulates of this theory are, (i) The tendency of an atom to take part in chemical combination is determined by the number of valence electrons. The valence electrons are the electrons in the outermost shell of the atom. (ii) The atoms combine by mutual sharing or by transfer of one or more electrons. In doing so, each combining atom acquires stable noble gas electronic configuration having 8 electrons in its outermost shell. This is called octet rule. (iii) The number of electrons which an atom loses, gains or mutually shares to attain noble gas configuration is called its valency. For example, Li, Be, B and C having respectively 1, 2, 3 and 4 electrons, have valence of 1, 2, 3 and 4. The elements N, O, F and Ne having 5, 6, 7 and 8 electrons in their outermost shell show common valence of 3, 2, 1 and 0. Thus, "the common valency of an element is either equal to the number of valence electrons, or it is equal to 8 minus the number of valence electrons."
LEWIS ELECTRON DOT SYMBOLS An American chemist, G.N. Lewis introduced simple notation to denote the valence electrons in an atom. These notations are called electron dot symbols or Lewis symbols, (or Lewis structures). According to this method,(a) The symbol of the element represents the nucleus along with all the inner electrons which do not take part in the bond formation.(b) The dots on the symbol represent the valence electrons. Thus, the number of dots represents the number of valence electrons. For example,(i) An atom of hydrogen contains one electron in its valence shell. So, its Lewis symbol is H. Here, H represents the nucleus of hydrogen and the dot (×) represents one valence electrons in an atom of hydrogen.(ii) The electronic configuration of chlorine is 2,8,7. Thus, there are seven valence electrons. The Lewis symbol of chlorine atom is
Recapitulation Exercise • Draw the Lewis symbols for the following elements. Na, Ca, B, Br, Xe, As, Ge • Why are Noble gases monoatomic? • Draw Lewis Symbols of O2- , Mg2+ ions. • Define Octet Rule.
Sub Topic 2: Types of Chemical Bonds The type of chemical bond developed between the two combining atoms depends upon the way these atoms acquire a stable noble gas configuration. Elements may combine through any one of the following ways to form stable compounds. • By the transfer of electrons from the atom of an element to the atom or atoms of another. This gives rise to an ionic (or electrovalent) bond. • By mutually sharing the electrons. This gives rise to a covalent bond. • By one-sided sharing of electrons. This gives rise to a coordinate bond.
Ionic Bond An IONIC BOND is an electrostatic interaction that holds together a positively charged ion (cation) and a negatively charged ion (anion). In an ionic bond, one atom loses an electron to another atom, forming a cation and anion, respectively. And, as everyone knows, opposites attract. Formation of NaCl
Sodium chloride results from ionic bonding. In table salt, for example, a valence electron from a sodium atom is transferred to a chlorine atom, forming Na+ and Cl-. Because the ions have opposite charges, they are attracted to each other. The loss of a valence electron and the attraction to the atom that took it happen simultaneously.
Properties of Ionic (or electrovalent) BondAn ionic or electrovalent bond has the following characteristics. : (i) An ionic bond is formed due to the coulombic attraction between the positively and negatively charged ions. (ii) An ionic bond is non-directional, i.e., the strength of interaction between two ions depend upon distance, but not on the direction. (iii) An ionic bond gets broken when the substance is dissolved in a polar solvent such as water, or when the substance is melted. Formation of Some Typical Ionic CompoundsThe formation of some ionic compounds is explained below. 1. Formation of magnesium chloride, (MgCl2).The electronic configuration of magnesium (At. No. 12) is 2,8,2. So, it has two electrons in its valence shell. The electronic configuration of chlorine (At. no. 17) is 2,8,7. So, it has seven valence electrons. In terms of the Lewis (electron dot) structures, one can write,
2. Formation of aluminium fluoride (AlF3). Factors Influencing the Formation of an Ionic Bond Formation of an ionic bond is favoured by, (i) Low ionisation enthalpy of the metallic element which forms the cation. (ii) Large electron gain enthalpy (electron affinity) of the non-metallic element which forms the anion. (iii) Large lattice energy, i.e., the smaller size and higher charge of the ions. Lattice Energy The strength of binding forces in solids is described by the term lattice enthalpy (DLH) (earlier the term lattice energy was used). The molar enthalpy change accompanying the complete separation of the constituent particles that composed of the solid (such as ions for ionic solids and molecules for molecular solids) under standard conditions is called lattice enthalpy (DLH°).
The lattice enthalpy is a positive quantity. For example, The enthalpy change for the reaction, under standard conditions. NaCl(s) Na+(g)+ Cl-(g) is the lattice enthalpy of NaCl(s), (DLH°(NaCl(s)) and that for the reaction, H2O(s) H2O(g) is the lattice enthalpy of the molecular solid ice. The lattice enthalpy of a molecular solid is the same as its standard enthalpy of sublimation. The lattice enthalpy of a metal is the same as its enthalpy of atomisation. The enthalpy change for the reaction, Na+(g) + Cl-(g) ® Na+Cl-(crystal) + lattice Enthalpy (U). is equal to -DLH (NaCl,s), i.e., heat equal to the lattice enthalpy is released during the formation of crystalline sodium chloride from gaseous ions. Lattice enthalpies (energies) are usually estimated from the thermochemical data using the Born-Haber cycle or by theoretical calculations. It depends upon the following factors: • Size of the ions. Smaller the size of the ions, lesser is the internuclear distance. Consequently the inter-ionic attractions will be high & the lattice enthalpy will also be large. • Charge on the ions.: Directly proportional to the magnitude of charge on the ions.
BORN-HABER CYCLE This thermochemical cycle was devised by Born and Haber in 1919. It relates the lattice energy of a crystalline substance to other thermochemical data. The Born-Haber cycle is the application of Hess's law to the enthalpy of formation of an ionic solid at 298 K.
Formation of crystalline sodium chloride form sodium metal and chlorine gas can be described by the reaction. Na(s) + ½ Cl2(g) ® NaCl (crystal) DrH = DfH° = - 411 kJ mol-1 (energy evolved) This overall reaction can be considered to proceed in a stepwise manner as follows The signs of the energy involved in each step follow the rule that energy evolved is negative and energy absorbed is positive. These steps are summarized in Fig. 6.3. From the Born-Haber cycle the value for any one of the steps can be calculated if data for all the other steps are known.
Recapitulation Exercise • Name the factors which favour the formation of an ionic bond. • Out of MgO and NaCl, which has higher lattice energy and why? • Draw the Born-Haber for a simple ionic solid such as MX
Sub Topic 3: Covalent Bonds A covalent bond is formed between two atoms (similar or dissimilar) by a mutual sharing of electrons. The shared pairs of electrons are counted towards the stability of both the participating atoms. A covalent bond is defined as the force of attraction arising due to mutual sharing of electrons between the two atoms. The combining atoms may share one, two or three pairs of electrons. When the two atoms combine by mutual sharing of electrons, then each of the atoms acquires stable configuration of the nearest noble gas. The compounds formed due to covalent bonding are called covalent compounds. CovalencyThe number of electrons which an atom contributes towards mutual sharing during the formation of a chemical bond is called its covalency in that compound. Thus, the covalency of hydrogen in H2 (H - H, H H) is one; that of oxygen in O2 is two (O = O, O O), and that of nitrogen in N2 is three (N º N, N N)
Characteristics of a Covalent Bond A covalent bond has the following characteristics. (i) Mode of formation. Covalent bonds are formed due to mutual sharing of one or more pairs of electrons. (ii) Directional character. Covalent bonds are directional in nature. This is because in a covalent bond, the shared pair of electrons remains localised in a definite space between the nuclei of the two atoms. This gives a directional character to the covalent bond. Single Covalent Bond A covalent bond formed by mutual sharing of one pair of electrons is called asingle covalent bond, or simply a single bond. A single covalent bond is represented by a small line (-) between the two atoms.
Formation of ammonia (NH3).The electronic configurations of nitrogen and hydrogen are N 1s2 2s2 2p3 or 2,5 H 1s1 or 1 Thus, each nitrogen atom requires three more electrons to acquire a stable noble gas configuration. On the other hand, each H-atom requires only one electron to achieve the stable helium configuration. This is done by mutually sharing three pairs of electrons between one nitrogen and three hydrogen atoms, as shown below. The unshared pair of electrons on the nitrogen atom (in ammonia molecule) is not involved in bond formation and is called a lone pair of electrons. Since, lone pair of electrons does not take part in bonding, hence it is also called non-bonding pair of electrons.
MULTIPLE COVALENT BONDS: The covalent bonds developed due to mutual sharing of more than one pairs of electrons are termed multiplecovalent bonds. These are, Double covalent bond.The bond formed between two atoms due to the sharing of two electron-pairs is called a double covalent bond or simply a double bond. It is denoted by two small horizontal lines (=) drawn between the two atoms, e.g., O = O, O = C = O etc. Triple covalent bond.Bond formed due to the sharing of three electron-pairs is called a triple covalent bond or simply a triple bond. Three small horizontal lines between the two atoms denote a triple bond, e.g., N º N, and H - C º C H (acetylene). Formation of Molecules Having Double & Triple BondsSome typical molecules having double & triple bonds are described below.
Oxygen molecules shares two electrons to make a double covalent bond. In a molecule of carbon, two atoms share three electrons -- a triple covalent bond. Formation of oxygen (O2) molecule. Formation of ethyne (C2H2) molecule.
Comparison Between Single, Double and Triple Covalent BondsSingle, double and triple covalent bonds differ from each other in the following ways. (i) Single bond is formed by the sharing of one electron pair, (two electrons), double bond is formed by the sharing of two electron pairs, (four electrons), whereas a triple bond involves sharing of three electron pairs, (six electrons). (ii) In a triple bond, six electrons attract the nuclei with greater force. This decreases the distance of separation between the two nuclei. In a double bond, four electrons attract the nuclei with a relatively lesser force, and in a single bond, two electrons hold the nuclei with a still lesser force. Therefore, the bond lengths follow the order, Triple bond length < Double bond length < Single bond length Since, a shorter bond means greater bond strength hence, the energy required to separate the bonded atoms (called bond energy) follows the order, Triple bond > Double bond > Single bond Formation of a covalent bond is favoured by, (i) High ionisation enthalpy (or energy) of the combining elements. (ii) Nearly equal electron gain enthalpies (or electron affinities) and equal electronegativities of the combining elements. (iii) High nuclear charge and small atomic size of the combining elements.
COORDINATE COVALENT BOND Coordinate bond is formed when the shared electron-pair is provided by one of the combining atoms. The atom which provides the electron-pair is termed as the donor atom, while the other atom which accepts it, is termed as the acceptor atom. The bond formed when one-sided sharing of electrons take place is called a coordinate bond. Such a bond is also known as dativebond. A coordinate bond is represented by an arrow (®) pointing towards the acceptor atom. Formation of Coordinate Bond During the Formation of a Molecule or Molecular ionThe formation of such a bond is illustrated through some examples given below. (i) Formation of ammonium (NH4+) ion.During the formation of ammonium ion, nitrogen is the donor atom, while H+ is the acceptor ion as shown below.
EXCEPTIONS TO THE OCTET RULE The octet rule is very useful for describing bonding in a large number of compounds. However, there are many exceptions to this rule. (i) Where duplet is formed.A hydrogen atom has only one electron in its valence shell. It needs one more electron to fill its valence shell. The completed shell has the electronic arrangement of the noble gas helium. In this case, therefore, an octet is not completed, but we still get a stable molecule. The electron dot structures of a few molecules containing hydrogen atom are shown below:
(ii) Where the octet remains incomplete.The elements of group 1,2 and 13 contain less than four electrons in their valence shell. These elements therefore, cannot achieve an octet by electron sharing. As a result, therefore these elements should not form covalent compounds. But, elements of these groups form some covalent compounds. Boron halides (BF3 and BCl3) are covalent compounds where the octet is incomplete (only six electrons surround the boron atom). These compounds are thus electron-deficient compounds. (iii) Where the octet is expanded.The elements belonging to groups 15, 16 and 17 have more than four electrons in their outermost shell. The elements of these groups form stable compounds in which there are more than eight electrons around the central atom. For example, PF5, PCl5 and SF6 are some typical compounds of this type.
The Cl and F atoms have 7 electrons in their valence shell. Therefore, they need one more electron each to attain the noble gas configuration. Phosphorus atom has five, and sulphur atom has six electrons in their valence shell. Then, the Lewis structures of PF5 PCl5, and SF6 are written as follows. Thus, P and S atoms have expanded their shells to accommodate more than eight electrons. Here again, octet rule is violated, but the compounds formed are stable.
Recapitulation Exercise • Give one example of a compound containing double bond & one containing a triple bond. • Describe a coordinate bond, giving one example. How does it differ from the Covalent bond? • How does bond multiplicity affect the bond length? • Explain the term Electrovalency & Covalency. • Write one main difference between an ionic and a covalent bond. • Where are the bonding electrons most likely to be found in a diatomic covalent molecule? • What happens to the valence electrons when a covalent bond is formed between two atoms?
Sub Topic 4: VSEPR Theory & Shapes of Molecules The VSEPR theory was proposed by R.J. Gillespie and R.S. Nyholmm in 1957. This theory was developed to predict the shapes of the molecules in which the atoms are bonded together with single bonds only. This theory is based on the repulsions between the electron-pairs in the valence-shell of the atoms in the molecule. The main postulates of the VSEPR theory are: (i) The geometry of a molecule is determined by the total number of electron pairs (bonding and non-bonding) around the central atom of the molecule. The shape of the molecule depends upon the orientation of these electron pairs in the space around the central atom. (ii) The electron pairs (shared, or lone pairs) around the central atom in a molecule tend to stay as far away from each other as possible so as to minimize the repulsion forces between them. (iii) The strength of repulsions between different electron pairs follows the order: Lone pair - Lone pair > Lone pair - Shared pair > Shared pair Shared pair The shared pairs of electrons are also called bond pairs of electrons. The presence of lone pair(s) of electrons on the central atom causes some distortions in the expected regular shape of the molecule.
Predicting the Shape of Molecules on the Basis of VSEPR TheoryAccording to the VSEPR theory, the geometry of a molecule is determined by the number of electron-pairs around the central atom. So, to use this theory for predicting the shapes of molecules just count the number of electron pairs (both, shared and lone pairs). The use of this theory in predicting the shapes of molecules is illustrated below by taking a typical molecule of the type ABn, where A is the central atom, B atoms are bonded to A by single electron pair bonds (single covalent bonds), and n is the number of B atoms bonded to one atom of A. For the sake of easier understanding we have divided molecules into various categories Shapes of the molecules having only the bond (shared) pairs of electrons(i) Molecules with two bond pairs.In a molecule having two bond pairs of electrons around its central atom, the bond pairs are located on the opposite sides (at an angle of 180°), of the central atom so that the repulsion between them is minimum. Such molecules are therefore linear. For example, in a molecule of the type AB2, in which the central atom A has two electron pairs, the two electron pairs are located on either side of A. Thus, the molecule AB2 takes a linear geometry. Some molecules which show linear geometry are, BeF2 (beryllium fluoride), BeCl2 (beryllium chloride), BeH2 (beryllium hydride), ZnCl2 (zinc chloride), and HgCl2 (mercuric chloride).
(ii) Molecules with three bond pairs.In a molecule having three bond pairs of electrons around its central atom, the electron pairs form an equilateral triangular arrangement around the central atom. Thus, the three bond pairs are at 120°C with respect of each other. Therefore, the molecules having three bond pairs around its central atom have trigonal planar (or triangular planar) shape. For example, in a molecule of the type AB3, the three bond pairs of electrons are located around A in a triangular arrangement. Thus, the molecule AB3 has a triangular planar geometry. Some molecules which show triangular planar geometry are; BCl3, BF3 etc. (iii) Molecules with four bond pairs (AB4 TYPE).
(iv) Molecules with five bond pairs.A molecule having five bond pairs around its central atom has a triangular bipyramidal shape. (v) Molecules with six bond pairs.The molecules of the type AB6 are octahedral. The molecule SF6 has an octahedral geometry.
Shapes of the molecules having bond pairs and lone pairs of electronsThe pair of electrons in the valence shell of an atom which is not involved in bonding is called lone pair of electrons. For example, the nitrogen atom in ammonia molecule has one lone pair of electrons; the oxygen atom in water molecule has two lone pairs of electrons. We describe a few examples of the molecules having one or more lone pairs of electrons. (i) Molecules having three bond pairs and one lone pair.A molecule of ÄB3 type having three bond pairs and one lone pair has a triangular pyramidal shape. Typical molecules of this type are NH3, NF3, PCl3, H3O+ etc. (ii) Molecules with two bond pairs and two lone pairs.The four electron pairs (two bond pairs + two lone pairs) are distributed tetrahedrally around the central atom as shown in Fig. 6.4, for a molecule . The two lone pairs on the central atom repel the bond pairs slightly inwards due to greater lone pair - bond pair repulsion. As a result, the bond angle in such a molecule is less than the tetrahedral value of 109°28¢.
These two representations of the H2O molecule show the electron density as a phantom shading (left) and contour lines (right.)Note how most of the negative charge is concentrated around the oxygen atom.
(iii) Molecules with four bond pairs and two lone pairs.The four bond pairs are distributed in a square planar distribution. The two lone pairs are in a direction at right angles to this plane. Thus, giving a square planar shape to such molecules. Examples include ICl4-, XeF4 and [Ni(CN)4]2-. (iv) Molecules with Five Bond Pairs & one lone pair:
(iv) Molecules with Four Bond Pairs & one lone pair, three bond pairs & two lone pairs and two bond pairs & three lone pairs respectively: