Chemistry I  © Ray Lovegrove; contributors: Reema Lodha, Antonio Codina, Madhusree Bathini, Ray Lovegrove, Rajesh Rathod et al. Source: PEOI

 

Chapter 4

Bonding

CHAPTER 4:

Bonding


During the chapter 3 we built up an idea of what atoms are like. We know that they are made up of tiny sub-atomic particles - electrons, protons and neutrons. We also understand that specific atoms of a given element will always have a given number of protons and electrons but may differ in terms of the number of neutrons that they have. Chemistry, for the main part is not concerned with what atoms are, but very concerned with what atoms do , in particular how individual atoms join together to form more complex substances. It is the atomic electrons that determine the way that atom combines with other atoms to form complex compounds. This will be the only topic of this chapter and it is essential that we understand it well, as it is the key to much of the chemistry that follows.

Atomic Structure and Bonding

In 1914 W Kossel found that by reacting elements of group 1A and 2A with elements of groups VIA & VIIA, produced salt like compounds. He concluded that the ions achieved a noble gas configuration and that electrostatic forces hold these ions.
In 1916, the American chemist Gilbert Newton Lewis proposed that chemical bonds are formed between atoms because electrons from the atoms interact with each other. Lewis had observed that many elements are most stable when they contain eight electrons in their valence shell. He suggested that atoms with fewer than eight valence electrons bond together to share electrons and complete their valence shells.
While some of Lewis' predictions have since been proven incorrect (he suggested that electrons occupy cube-shaped orbitals), his work established the basis of what is known today about chemical bonding.
In 1919 Irving Langmuir extended Lewis's ideas and introduces concept of the covalent or homopolar bonding. The success of Schrodinger’s quantum wave theory in 1926, dealing with the atomic structure, directed the entire attempt to understand the covalent bond towards this new approach. W. Heitler and F. London threw new light on the nature of covalent bond. They introduced number of concepts that are now of general use in chemical theory. These concepts are used along with the Lewis’s theory to give more insight into chemical bonding.

Filling the orbitals Hunds Rule
The atom is a central, positively charged nucleus surrounded by extra-nuclear electrons. These extra-nuclear electrons are in such a number as to achieve electrical neutrality and are grouped in several shells - K, L, M, N... These shells are divided into subshells - orbitals called: s, p, d, f.. and the electrons occupy each of these orbitals (subshells) in pairs. The number of electrons that a shell can contain increases with the distance of the shell to the nucleus. The number of electrons in a n is the quantum shell number, 3ii i.e for K (first shell n is 1) = 2, L = 8, M = 18, N = 32 etc. Fig. 2 (Electrons in green dots).

Sodium (Na) is a alkali metal (group IA element) with atomic number 11 in Periodic table Elements. It has 11 protons in the nucleus and to achieve electrical neutrality it has 11 electrons in the K, L, M shells. These electrons are arranged as, 2 (K shell), 8 (L shell), 1 (M shell) i.e. 2. 8. 1 (11 electrons) Fig. 1. Calcium (Ca) is a alkali earth metal (group II element) with atomic number 20. It has 20 protons in the nucleus and to achieve electrical neutrality it has 20 electrons in the K, L, M, N shells. These electrons are arranged as, 2 (K shell), 8 (L shell), 8 (M shell), 2 (N shell) i.e. 2. 8. 8. 2 (20 electrons) Fig. 3.

We already mentioned that these shells have sub-shells or more accurately termed orbitals; s, p, d or f. These orbitals are occupied by paired electrons in an opposite spin. The s orbital can occupy 2 such electrons (e), p orbital 6e, d orbital 10e, f orbital 14e. And the orbital electrons (in superscript) represented as follows:- 1s2 (K shell); 2s 2, 2p6 (L shell); 3s2, 3p6, 3d10 (M shell); 4s2, 4p6, 4d10, 4f14 (N shell). The order in which the orbitals s, p, d, f... become filled with electrons, is governed by their energy level, i.e. lower energy orbital filled first;

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d....
––––––––––––––––––––––––––––––––>
Increasing orbital energies

The electronic configuration of Sodium - Na; 1s2, 2s2, 2p6, 3s1 and for Calcium - Ca; 1s2 , 2s 2, 2p6, 3s2, 3p6, 4s2. Electronic Configuration 3ii 3vi

Sodium
electrons 2.8.1

Electrons in KLMN Shells[P+N Nucleus]

Calcium
electrons 2.8.8.2

Fig. 1

Fig. 2

Fig. 3


When the required number of electrons (equal to the Atomic Number) has been accounted for, the electrons which occupy the outer quantum shell (valency shell) are regarded as the valence electrons. The valence electrons are most exposed to another atoms and will participate in the formation of chemical bonds.

The octet rule
When the elements combine to form chemical compounds, the electrons in the outermost shell or valency shell may be transferred from one atom to another or there may be a mutual sharing of the electrons i.e. atoms are held together by the exchange or sharing of electrons, to achieve a stable configuration or ‘octet’ (G. Lewis, 1916). Bonds are formed by sharing pairs of electrons between two nuclei, so that, each achieves a noble gas (Group VIII/O) configuration. Noble gases have complete filled orbitals (ns2 np6) that confer extra stability. Atoms are therefore considered to combine together in such a manner as to achieve this stable Octet configuration (octet rule). Atoms achieve this stable state chemically in three ways: Electrovalent (Ionic) Bonding, Covalent Bonding, and Dative (co-ordinate, semi-polar, semi-ionic) Covalent Bonding.

Valency and Bonding
In 1914 W Kossel found that by reacting elements of group 1A and 2A with elements of groups VIA & VIIA, produced salt like compounds. He concluded that the ions achieved a noble gas configuration and that these ions are held by electrostatic forces.
In Electrovalent (Ionic) bonding the valency or the combining power of an atom is equal to the number of charges on Cation (positive ion) or Anion (negative ion). Monovalent include; Li, Na, K, Rb, Cs, F, Cl, Br, I and Divalent; Mg, Ca, Sr, Ba, O, S, Se, Te.

Six Periodic
elements and the electrons in KLM shells

Table showing six periodic elements and the electron(s) in outer valence KLM shell Fig. 4

Positive Valency
In the Group I alkali metal atom, loss of an electron destroys the electrical neutrality (free state) of the atom and it becomes a positive ion (+1 valency). The element Sodium (Na) has 11 protons and 11 electrons in its free state. It has one electron in the outer M shell, that it loses to obtain noble gas configuration and become stable. On losing the electron, it will have + 1 charge (Cation) and the valency is +1 Fig. 5

Negative Valency
In the Group VIIA halogen atom there is one less electron (in outermost shell) than the noble gas (Group VIII/O) structure, so the capture of an electron by a halogen atom will give the eight electrons of the noble gas octet structure. As before, the electrical neutrality (free state) is destroyed, and the atom became a negative ion (Anion). The element Chlorine (Cl) has 17 protons and 17 electrons in free state. It has 7 electrons in the outer M shell, one less, to fit the stable octet configuration. On gaining the electron, it will have a –1 charge (Anion) and the valency will be –1 Fig. 6

positive valency

negative valency

Fig. 5

Fig. 6


In 1916 G Lewis proposed that bonds formed by sharing pairs of electrons between two nuclei so that each nuclei achieves a noble gas configuration. Langmuir extended Lewis's ideas - concept of the covalent or homopolar bond. In covalent bonding the valency of an atom is equal to the number of electron pairs shared with the other nuclei. Divalent; O, S, Se, Te, Trivalent; B, Al group, N, P group, Tetravalent; C, Si, Ge, Sn, Pb, Pentavalent; N, P, As group.

Electrovalent (Ionic) Bonding
The electrovalency of an element is the number of electrons its atom must gain or lose to attain a stable electron structure. In ionic (electrovalent) bonding, electrons are completely transferred from one atom to another. Ionic bonds are a type of electrostatic bond between two atoms and are formed from the mutual attraction of ions. Atoms that have charges are called ions. Ionic bonds can be explained using the theory of electron orbitals. Every atom has a number of electron orbitals equal to or greater than one. Each orbital has a maximum electron limit, after which a new orbital is created. The number of orbitals in the elements ranges from one for hydrogen or helium, to six, seven, or eight for the larger molecules with atomic numbers greater than uranium. Atoms are "most stable" when their outermost orbitals are filled. As the atoms approach each other, each nucleus begins to attract the electron held by the other nucleus. Eventually, the electron clouds overlap and fuse into one "molecular orbital". Like an atomic orbital, a molecular orbital is most stable when filled by a pair of electrons. This shared orbital acts as a "chemical bond" between the two atoms. In the process of either losing or gaining negatively charged electrons, the reacting atoms form ions. The oppositely charged ions are attracted to each other by electrostatic forces, which are the basis of the ionic bond.

Typical participants in an ionic bond are a metal and a non-metal, such as sodium and chlorine. We revisit our example of the reaction of sodium with chlorine that creates sodium chloride, or common table salt.
We have the positively charged Sodium ion and a negatively charged Chlorine ion. These attract each other because opposite charges attract.

Elements like the Alkali metals (Group IA) or halogens (Group VIIA) whose atoms lose or gain only one electron, have a valency of 1. As it is only the electrons in the outermost shell which are affected, A single electrovalent bond is formed by the sharing of two electrons. Sodium (Na) reacts energetically with Chlorine to form the very stable salt Sodium Chloride.The essential chemical action is the capture of an electron from each sodium atom by each chlorine atom. After they have reacted the atoms are now oppositely charged ions in close association, and powerful electrostatic forces of attraction hold them in contact Fig. 7. Ions combined together this way are said to be linked by an Ionic bond, and compounds formed from this type of binding are called Electrovalent or Ionic. A crystal of NaCl is a closed packed system.

sodium(Na) Group
I-Alkali Metal

Magnesium (Mg) Group
II-Alkali Earth Metal

Fig. 7

Fig. 8


The Alkali Earth metals in Group IIA have an electrovalency of 2, because two electrons must be removed to expose the noble gas structure in the underlying shell. Example shown in Fig. 8 is Magnesium (Mg).
In order to maintain overall electrical neutrality (free state), doubly charged ions are always associated with two single charged ions of opposite sign or another double charged ion of opposite sign. Other e.g. MgO (Mg2+ + O2- ).


Properties of Ionic or Electrovalent Compounds.

  1. At ordinary temperatures they invariably exist as crystalline solids although individual particles such as Na+ Cl- may exist in the vapour state. The crystal consists of an ‘infinite assembly’ of ions, that is, a large and indefinite number of oppositely charged ions joined together in a regular manner.

  2. Electrovalent compounds have high melting points and boiling points owing to the strong attraction between the oppositely charged ions.

  3. In the fused state they are good conductors of electricity and are electrolysed by a current. This is explained by the breakdown of the crystal into ‘free’ ions, which are attracted to, and discharged at, the electrode of opposite sign.


  4. They often dissolve readily in water, but are only sparingly soluble in organic solvents like ether and benzene.


  5. When oppositely charged ions react together in aqueous solution they do so almost instantaneously.

Covalent Bonding
When there is sharing of electrons between atoms, the bonds are called covalent or homopolar bonds. Valency of the atom is the number of electron pairs shared with the other nuclei. Transferring electrons is not the only way that atoms can combine, covalent bonding occurs when two (or more) elements share electrons.

Covalent bonding occurs because the atoms in the compound have a similar tendency for electrons (generally to gain electrons). This most commonly occurs when two nonmetals bond together. Because both of the nonmetals will want to gain electrons, the elements involved will share electrons in an effort to fill their valence shells.

According to modern atomic theory, atoms have electrons circling them in shells called orbitals. Each orbital has a maximum number of electrons, and each atom "wants" to max out its electrons in each orbital. The bonds are formed between twp atoms each of its electron orbitals seeking a similar number of other electrons to max out their electron shells. When the atoms are brought together, their electron shells intermingle, and create something called "molecular orbitals," where electrons wander freely between both atoms and orbit the nuclei of both.

Depending on the number of shared electron pairs, a covalent bond is characterized as a single bond, a double bond, triple bond, etc. Some of the metals with the highest melting points, molybdenum and rhenium, have quadruple bonds. Quintuple covalent bonds and sextuple covalent bonds are quite rare, and we have good reason to believe that nothing on the periodic table can go beyond a sextuple bond.

Single Covalent Bond
Most common type of chemical bonding is single covalent bonding, where one pair of valence electrons is shared by the two atoms.


Hydrogen molecule [H2] Fig. 8a

A good example of a covalent bond is that which occurs between two hydrogen atoms Fig. 8a. Atoms of hydrogen (H) have one valence electron in their first electron shell. Since the capacity of this shell is two electrons, each hydrogen atom will "want" to pick up a second electron. In an effort to pick up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H2. Because the hydrogen compound is a combination of equally matched atoms, the atoms will share each other's single electron, forming one covalent bond. In this way, both atoms share the stability of a full valence shell.

Methane
Fig 8b

In methane molecule (CH4) Fig. 8b, the Hydrogen (H) requires one electron to form the helium electron (duplet) structure and carbon needs four more electron to form the neon (octet) electron structure. Therefore by sharing the electrons of four hydrogen atoms with each electron of the carbon atom, all the atoms gain the noble gas (octet) electron structure. A pair of electrons shared in this way constitutes a covalent bond.

Multiple Covalent Bonds
For every pair of electrons shared between two atoms, a single covalent bond is formed. Some atoms can share multiple pairs of electrons, forming multiple covalent bonds.


Oxygen molecule [O2] Fig. 9

For example Fig. 9, oxygen (which has six valence electrons) needs two electrons to complete its valence shell. When two oxygen atoms form the compound O2, they share two pairs of electrons, forming two covalent bonds.

In Fig. 10* only the electrons in the outer valency shells are shown in these examples since electrons in the inner completely filled shell, are not involved in bond formation. (The outer shell electrons in the atoms can also be reprensented by dots and crosses instead of colours). These examples shows the formation of single (Cl2), double (O2), triple (N2) covalent bonds. Covalent linkages of the above type are non-polar because each atom contributes the same number of electrons to the shared electron pairs and no separation of charges arises (other examples; NH3, BF3 ).
The multiple bond in N2 molecule, is formed by, each of the two atoms in the molecule of nitrogen completing its octet by sharing three of the five electrons belonging to the other atom. Two electrons with opposite spins constitute the ordinary covalent bond, and as the six shared electrons in the nitrogen molecule form three such pairs the result is three covalent bonds joining two atoms. Acetylene (C2H2) also has triple bonds. Double bonds are formed in similar way; e.gs. ethylene (C2H4), carbon dioxide (CO2). Double bonds are more common than triple bonds.

Single, Double and
Triple covalent bonds
Fig. 10

The main contribution that Lewis made to structural theory was the suggestion that a pair of electrons could be shared between two atoms to form a bond while still contributing to the stable octet of each (or duplet in the case of the H atom).

A Shorthand Method for Drawing Covalent Bonds
The problem with this "solar system" diagram of the atoms is that sharing two pairs of electrons gets even more confusing in molecules having triple or quadruple covalent bonding. Thus a method to simplify a diagram of the molecule was devised. It only shows the valence electrons as dots. It is called the electron dot notation or Lewis dot notation.
Lewis dot structure is shorthand to represent the valence electrons of an atom. Valence electrons are those that are in the outer orbit or shell of an atom. The structures are written as the element symbol surrounded by dots that represent the valence electrons.Fig 10a

Lewis dot notation
Fig.10a

Lewis structures for
Hydrogen and Oxygen

Lewis structures for H2 and O2 Fig.10b

Polar and Nonpolar Covalent Bonding
There are, in fact, two subtypes of covalent bonds.

Whenever two atoms of the same element bond together, a nonpolar bond is formed. The H2 molecule is a good example of the nonpolar bond. Because both atoms in the H2 molecule have an equal attraction (or affinity) for electrons, the bonding electrons are equally shared by the two atoms and a nonpolar covalent bond is formed

A polar bond is formed when electrons are unequally shared between two atoms. Polar covalent bonding occurs because one atom has a stronger affinity for electrons than the other (yet not enough to pull the electrons away completely and form an ion). In a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons. A good example of a polar covalent bond is the hydrogen-oxygen bond in the water molecule.


Water molecule [H2O] Fig. 10c

Water molecules contain two hydrogen atoms bonded to one oxygen atom. Oxygen, with six valence electrons, needs two additional electrons to complete its valence shell. Each hydrogen contains one electron. Thus oxygen shares the electrons from two hydrogen atoms to complete its own valence shell, and in return shares two of its own electrons with each hydrogen, completing the H valence shells. Fig. 10c

Because the valence electrons in the water molecule spend more time around the oxygen atom than the hydrogen atoms, the oxygen end of the molecule develops a partial negative charge (because of the negative charge on the electrons). For the same reason, the hydrogen end of the molecule develops a partial positive charge. Ions are not formed; however, the molecule develops a partial electrical charge across it called dipole Fig. 16.


Properties of Covalent Compounds.

  1. They are Low melting solids or liquids (due to the absence of strong electrostatic intermolecular forces).
  2. Formed between nonmetals.
  3. Covalent bonds tend to exist as true molecules.
  4. They do not conduct electricity.

  5. Covalent molecules tend to exist as liquids or gases at room temperature. H2O (water) is a liquid at room temperature.
  6. Sparingly soluble if not insoluble, in water.

Dative (Co-ordinate) Covalent Bonding
Molecules with a lone pair of electrons like those of Methanol (CH3OH), Ammonia (NH3 ), Oxygen (O2) or Water (H2O), frequently behave as electron donors and use the lone pair to form a covalent bond with an electron acceptor. Molecules which behave as electron acceptor have available vacant orbitals and gain extra stability from the additional electrons. Boron fluoride is an electron acceptor and forms a stable additional compound with Methanol Fig. 11.

Methanol + Boron
Fluoride

Fig. 11

As a result of an atom losing a share of two electrons, it is left with a partial positive charge and the electron acceptor that gains the extra share of electrons, acquires an equal partial negative charge. So that associated with the co-ordinate bond, there is also a weak ionic bond and the co-ordinate bond, could equally be represented as A б+—>B б- (the symbol implies a partial charge on the atom).
Because of this it has been called a Semi-polar bond. The term dative bond is also used. Co-ordinate bond is represented by an arrow sign with or without partial charges.
Another example is Ammonium Chloride (NH4Cl), which contains electrovalent and simple covalent linkages as well. Covalently bound to the N atom in the NH4+ ion in this way, the H+ ion from hydrogen chloride is indistinguishable from the other H atoms, and the positive charge is spread symmetrically over the whole structure giving it the characteristics of a unipositive ion. This then combines ionically with negative Cl- ion in the usual way, ultimately producing an ionic solid Fig. 12.

Ammonium Chloride

Fig. 12

The properties of a particular salt are conditioned by the size and charge of the cation. The alkali metals form unipositive ions, which have large ionic radii. The NH4 radical has a single positive charge and a large radius, this explains why there are similarities between NH4salts and the salts of alkali metals.

Metallic Bonding
We can use the idea of ionic and covalent bonding to explain all of the compounds formed between metals and non-metals, non-metals and other non-metals and elements in which more than one atom is usually present ( H2, O2, N2 etc.). However, most of the elements in the periodic table are metals and we need to look at how these bond together. Metal atoms, as you will remember, all have very few electrons in their outer shells. When metal atoms get together, the outer electrons (depending on the element concerned, this will be one, two or three) are given away, changing the atoms to cations. All the electrons form a negative charge 'cloud' or 'sea' surrounding the cations. The negative sea of electrons acts like a glue which holds the cations together.

All the properties of metals are a direct result of metallic bonding:


We will look more closely at the structure of metals in chapter six.

Structures
It would be unwise to consider the bonding between atoms as the only way that we can understand them. Chemists use the word bonding to think about what is happening on a very small scale. If we wish to think on a bigger scale that considers how substances behave, chemists use the word Structure. Table below shows Structure, type of bonding and examples.

STRUCTURE
BONDING
EXAMPLES
Simple molecule
Covalent, sometimes Dative covalent
Water, methane, hydrogen gas.
Macromolecule
(Giant covalent lattice)
Covalent
Diamond, graphite and many silicate rocks.
Ionic lattice
Ionic
Sodium chloride, magnesium bromide.
Metallic lattice
Metallic
Sodium, iron.

Table showing Structure, type of bonding and examples.

Diamond
The structure of diamond is an excellent example of how closely the properties of a substance are related to its structure. Diamond is made up of carbon atoms, each forming four strong covalent bonds with its neighbours. As a covalent bond is strong, four covalent bonds (in a three dimensional model at 109° to one another) make for a very strong structure indeed. Diamond is what chemists call a 'giant' covalent structure, namely that a single diamond is one structure, made up perhaps of millions of carbon atoms. The unique optical properties of diamonds are due to the fact that the atoms are arranged in a very clear pattern. Diamonds, apart from being 'a girl's best friend,' have many uses where a hard and strong substance is vital, such as a dentist's drill. It is possible to make artificial diamonds - albeit with some difficulty - but these are reserved for industrial use as their appearance makes them unacceptable for jewellery Fig. 13.

Graphite
Like diamond, graphite is a giant covalent molecule composed entirely of carbon atoms. Unlike diamond, graphite is not particularly strong and does not transmit light; in fact it is a rather dull, grey solid. In the graphite structure, carbon atoms are joined to each other by three, strong covalent bonds, forming a layered structure Fig. 14. If you think of the carbon atom you will remember that it has four electrons in its outer shell, the fourth electron from each atom being free to move around. Electrons that move will of course carry a charge, and that is why graphite is unusual among non-metals (but not unique) in that it conducts electricity. The free electrons in graphite behave in a similar way to the electrons in a metal (see above).


Diamond

Graphite

Fig. 13

Fig. 14


Intermolecular Bonding
So far, we have restricted our discussions to the bonding that takes place between atoms - ionic, covalent and metallic bonding. There are a number of weaker intermolecular attractions that occur between molecules. These are called "intermolecular bonds" and can be defined as a weak attractive force between molecules, caused by a small difference in electrostatic charge. While these bonds are very easily broken, they do have a great influence on the physical properties of a substance, such as its melting and boiling point.
The most common of these intermolecular attractions are called VAN DER WAALS forces and exist when small areas of negative (δ-) and small areas of positive charge (δ+) form on the surface of the atom for a very short time. Even inert atoms like helium have momentary small areas of charge on their surface due to the movement of electrons; this very small attractive force enables helium to be liquified at very low temperatures (-269°C or 4K). An element that had no attractive forces between atoms would remain a gas even at absolute zero (-273°C or 0K) - such particles do not exist.

Note. The temperatures given as 4K and 0K are stated in the Kelvin scale - this is a scale in which absolute zero is 0K and, as you can not have a temperature below this point, there can be no negative values. As the Kelvin is identical to a degree Celsius, you simply add 273 to the Celsius temperature to give the absolute temperature ( for instance 27°C is 27 + 273 = 300K) You should note that no degree symbol " o " is used when the temprature is stated using the Kelvin scale.

A very important form of intermolecular bonding is the HYDROGEN BOND of which we will have more to say in later chapters. Hydrogen bonding occurs in molecules where hydrogen is bonded covalently to one of a few other elements, most importantly oxygen and nitrogen. These elements have a much stronger pull on the electrons in a bond than the hydrogen atom, and are said to have a greater ELECTRONEGATIVITY 3vii. Electronegativity is a measure of how good atoms are at attracting electrons; a numerical value is given to each element which is nothing more than an index, a way of comparing it with the other elements. Electronegativity values have no units (you generally don't need to refer to the actual values for electronegativity. Electronegativity increases across the period and decreases down the group e.g the value increases from Li (1.0) to F (4.0) and decreases from Li (1.0) to Cs (0.7) Fr (0.7). Metals have lower values and non metals - especially those in groups VA (15) - VIA (16) and VIIA (17) of the periodic table - have high values). If a covalent bond is formed between two atoms of differing electronegativity, the electrons are unevenly distributed and the bond is described as POLAR. A bond between two atoms of similar electronegativity is not going to have an unequal distribution of charge and the bond is called NON-POLAR. In the diagram below Fig. 15, all the bonds are polar except the H-C bond.

All the bonds are
polar except the H-C bond

water

Fig. 15

Fig. 16

The way in which charges on a water molecule are distributed is important, as it shows how hydrogen bonds are formed. The hydrogen atoms have their electrons involved in covalent bonds with oxygen. This leaves the hydrogen nuclei exposed, giving a small positive area of charge to this area of the molecule (δ+). The oxygen atom of the water molecule has all of the electrons in the molecule gathered around it, giving this part of the molecule a small negative charge (δ-).
In the diagram Fig. 16 we can see how hydrogen bonds form between the water molecules ( hydrogen bonds are shown by a dotted line ----------). It is hydrogen bonding that accounts for the very unusual behaviour of water (see Chapter 10).

Molecular orbital and Bonding
As the atoms are brought together, the 1s atomic orbital of hydrogen overlap and coalesce to form molecular orbital, the contour of which defines the region, which both electrons (provided that they are of opposite spin) can be considered to occupy jointly, Figure below.

The molecular orbital Hydrogen is ovoid and symmetrical about line joining the two nuclei and paired electrons constitute single bond between nuclei and is defined as a sigma bond (σ). The hydrogen molecule is envisaged as two positive nuclei surrounded by electron cloud. The Hydrogen molecule may be described as two positive nuclei surrounded by an electron cloud.
In chlorine molecule half-filled 3p orbital are involved. Stability of the Hydrogen molecule is due to the fact that electrons are within larger region of the molecular orbital, although greatest electron density is found in the region between the two nuclei i.e. shared electron pair.
In dative co-ordinate covalent bonding, the bond is formed by the overlapping of an atomic orbital occupied by an electron pair with a suitable vacant orbital.
The electronic configuration of Carbon [C] is 1s2, 2s2, 2px1, 2py1 (K 2, L 4), in which px & py are half filled, and C is divalent. It is concluded that the quadrivalent nature of carbon is due to one of the 2s electrons being promoted to the vacant 2pz orbital thus giving four half-filled orbital.

Hybridisation
When an atom is in an ‘excited’ state through absorption of energy, unpairing of its electrons may occur. If there is an unoccupied orbital in a higher energy sub-level one of the unpaired electrons is ‘promoted’ to the vacancy and both electrons become available for bond formation. Beryllium atom in its ground state has an electron configuration 1s2 2s2 . When beryllium combines with chlorine the 2s electrons become unpaired and one is ‘promoted’ to one of the three empty orbitals in the 2p sublevel. One might expect that the beryllium atom would combine with two chlorine atoms by means of an s-p bond and p-p bond. This would result in two bonds of different strengths, a p-p bond being stronger than an s-p bond because of greater overlapping of the charge clouds. Actually the two bonds in Beryllium Chloride (Be Cl2) are equal in strength. This is because of a process known as hybridisation of atomic orbitals. The wave systems of the two-beryllium electrons interact to give two similar orbitals, with axes at 180 degrees to each other. The beryllium chloride molecule is thus linear (Cl—Be—Cl). Since the new orbitals are derived from one s orbital and one p orbital we describe them as ‘sp hybrid’ orbitals. The advantage of hybridisation of atomic orbitals is that the hybrid orbitals are concentrated in particular directions, so that there is greater overlapping of charge clouds, producing stronger bonds. . Similar sp hybridisation of orbitals occurs in the mercury atom when it combines with two chlorine atoms. Mercury (II) Chloride also has a linear molecule (Cl—Hg—Cl).


Fig. 17

Boron atom has electronic confirguration 1s2 2s2 2p1 (one more electron than beryllium atom). Two of the three p orbitals in the second quantum shell are unoccupied, one of the empty p orbitals, giving two p electrons. Hybridisation occurs between the remaining s orbital and the two p orbitals, and three equivalent sp2 hybrid orbitals are produced. The axes of these lie in the same plane at 120 degrees to each other Fig. 17.

The Carbon atom again has one more electron than the boron atom. Its electron configuration is 1s2, 2s2, 2px1, 2py1 (K 2, L 4), the 2px and 2py signify that two different p orbitals in the second quantum are occupied by single electrons. ‘Excitation’ of the carbon atom results in the carbon atom results in one of the 2s electrons being 'promoted' to the empty 2pz orbital, so that now all three p orbitals contain one electron. Hybridization takes place between the s orbital and the three p orbitals, giving four equal sp3 hybrid orbitals Fig. 18. These are directed outwards from the carbon atom so that they form the tetrahedral angle of 109 degree 28’minutes with each other. Thus a molecule of methane or carbon tetrachloride has the shape of a regular tetrahedron, the carbon atom being at the center and the hydrogen or chlorine atoms at the four corners Fig.19 e.g.


Fig. 18


Fig. 19

The four Hydrogen atoms in methane CH4 is considered as being bonded by the overlap of Hydrogen 1s orbitals with the half-filled orbitals of the Carbon atoms because of the s orbital being non-directional and p orbitals are directional, the four bonds (right angles to each other) in CH4would not be equivalent in direction. Also abnormality arises between experiment and theory of these four bonds, as they would neither be equivalent in direction (due to s orbital being non-directional & p orbital are at right angles to each other) nor be of equivalent strength (as s orbitals have lower energy than p orbitals). It has been found experimentally that 4 C––– H bonds in CH4 to be equivalent in direction and reactivity and the orbitals that constitute this must also be equivalent.
The abnormality is solved by modifying the direction and energies of the orbitals by mathematical solution i.e. hybridisation of one s orbital and the three p orbitals (sp3 hybridisation) leads to replacement of these orbitals with four new hybrid orbitals of equal energy which are directed towards corners of the bonds in the methane molecule, result from the overlap of equivalent hybrid in orbit with hydrogen 1s orbital to form experimentally verified tetrahedral molecule [image methane]. Other modes of hybridisation of C atomic orbitals are possible.

In formation of the ethene molecule (C2H4) H2C: :CH2 which contains a ‘double’ bond, hybridisation of atomic orbitals of the ‘exited’ carbon atom takes place differently. ‘After promotion’ of one of the 2s electrons to the 2p sub-level the remaining s orbital combines with only two of the p orbitals (sp2 hybridisation), so that each carbon atom has three sp2 hybrid orbitals and one unchanged p orbital Fig. 20. As with the boron atom, the axes of the sp2 hybrid orbitals lie in the same plane and are at 120 degrees at each other. Two of the hybrid orbitals are used in bonding with two hydrogen atoms, and the third in establishing a single bond with the other carbon atom. The axis of the remaining p orbital is at right angles to the axis of the hybrid orbitals, and the second bond between the carbon atoms is formed by overlapping of the two p orbitals or charge clouds.


Fig. 20




Fig. 21

Maximum overlapping will occur if the orbitals have the same symmetry with respect to the C — C axis, that is, if all six atomic nuclei lie in the same plane. Overlapping then takes place both above and below the plane.
Two types of bonds involved in hybridisation.
1] A Sigma bond is formed due to maximum overlapping between two nuclei (internuclear axis)- called ‘collinear overlapping’ i.e. covalent bond formed by collinear overlapping of atomic orbitals is called a Sigma (σ) bond. Examples of sigma bonds with collinear overlapping are H—H, O—H, C—H in methane (all sigma bonds).
2] A bond formed by ‘collateral’ overlapping of p orbitals is known as a pi (π) bond. Stronger of two bonds is Sigma bond.
Ethene cannot spin or twist because of the pi bond Fig. 21.

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