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Chapter 3

Bonding

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Electron Gain Enthalpy (Δ egH)

When an electron is added to a neutral gaseous atom (X) to convert it into a negative ion , the enthalpy change accompanying the process is defined as Electron Gain Enthalpy ( ΔegH). Electron gain enthalpy provides a measure of ease with which an atom adds an electron to form anion as represented by equation :

                    X(g) + e X ; Δ H = ΔegH ………. (1)

Depending upon the element, the process of adding an electron to the atom can be either endothermic or exothermic. For many elements energy is released when an electron is added to the atom and the electron gain enthalpy is negative. For example, group 17 elements (the halogens) have very high negative electron gain enthalpies because they can attain stable noble gas electronic configurations by picking up an electron. On the other hand noble gases have large positive electron gain enthalpies because the electron has to enter the next higher principal level leading to a very unstable electronic configuration. It may be noted that the electron gain enthalpies have large negative values towards the upward right of the periodic table preceeding the noble gases.

The variation in electron gain enthalpies of elements is less systematic than ionisation enthalpies.

Electron Gain Enthalpies (kJ/mol) of Some Main Group Elements

H

- 73

 

 

He

+48

Li

- 60

Be

+66

B

- 83

C

- 122

N

+31

O

- 141

F

- 200

Ne

+116

Na

- 53

Mg

Al

- 50

Si

- 119

P

- 74

S

- 200

Cl

- 349

Ar

+96

K

- 48

Ca

Ga

- 36

Ge

- 116

As

- 77

Se

- 195

Br

- 325

Kr

+96

Rb

- 47

Sr

I n

- 29

Sn

- 120

Sb

- 101

Te

- 190

I

- 295

Xe

+77

Cs

- 46

Ba

T l

- 30

Pb

- 101

Bi

- 110

Po

- 174

At

- 270

Rn

+68

Successive Electron affinity

Like the second and higher ionisation energies , the second and higher electron affinities are also possible. After addition of one electron, the atom becomes negatively charged and the second electron is added to a negatively charged ion. The addition of second electron is opposed by the coulombic force of repulsion and energy has to be supplied for the addition of the second electron. If an atom has spontaneous tendency, i.e., a positive tendency , to gain electron, then conventionally , its electron gain enthalpy is said to be negative and if the atom is reluctant to gain an electron, i.e., it has a negative tendency to gain an electron and is forced to accept it, its electron gain enthalpy is positive. Thus in the case of oxygen, the first electron gain enthalpy is negative since 141 k J is released when one mole of oxygen atoms get converted to O ions. In other words oxygen atom has positive tendency to accept electron. However, the second electron gain enthalpy is positive since 770 k J of energy has to be supplied to convert 1 mol of O ions to O2 ions.

Similarly in the case of sulphur, while first electron gain enthalpy is negative since 200 k J of energy is released when 1 mole of S atoms get converted to S ions and second electron gain enthalpy is positive since 590 kJ of energy has to be supplied to convert 1 mole of S - ions to S2 ions.

Thermodynamically , the energy released is given a negative sign and energy absorbed is given a positive sign . Accordingly, when a species has a positive electron affinity , ΔH , accompanying the addition of an electron to the species, is negative and if it has a negative electron affinity , ΔH is positive. Thus for the reaction Cl + e Cl, while electron affinity is positive , while electron gain enthalpy is negative. In various calculations involving ΔH, the value of electron affinity of chlorine would be taken as −349 k J mol−1and not +349 k J mol−1.

Factors on which Electron Gain Enthalpy depends

The important factors upon which electron gain enthalpy depends are briefly discussed.

  1. Atomic size : As the size of an atom increases , the distance between its nucleus and the incoming electron also increases. Consequently, the incoming electron experiences less attraction towards the nucleus of the atom. Therefore , electron gain enthalpy becomes less negative down the group.
  2. Nuclear charge : With increase in the nuclear charge, force of attraction between the nucleus and and incoming electron increases and so is the value of electron gain enthalpy. Thus, the electron gain enthalpy becomes more negative with increase in nuclear charge.
  3. Symmetry of electronic configuration : The symmetry of electronic configuration has very important role to play. The atoms with symmetrical configuration (having filled and half-filled orbitals in the same sub-shell) do not have any urge to take up extra electrons because their configuration will become unsymmetrical or less stable. In case these are made to accept electrons, energy will be needed and electron gain enthalpy will be positive. For example , noble gas elements have positive electron gain enthalpies.

Variation of Electron gain enthalpy

Along a period

As a general rule, electron gain enthalpy becomes more negative with increase in atomic number across a period. The effective nuclear charge increases as we go from left to right across a period and consequently it will be easier to add an electron to smaller atom since the added electron on an average would be closer to the positively charged nucleus.

The trends in electron gain enthalpy values within a period are irregular indicating that atomic size is not the only criterion for determining electron gain enthalpy. Thus electron gain enthalpy of Be is positive (+66 k J mol−1 ) is while that of Li is negative (−60 k J mol−1). Similarly , electron gain enthalpy of nitrogen is positive(+31k J mol−1) while that of oxygen although atomic size of oxygen is less is negative (−141 k J mol−1 ) .

Along a group

We should also expect electron gain enthalpy to become less negative as we go down a group because the size of the atom increases and the added electron would be farther from the nucleus. This is generally the case (Table above). However, electron gain enthalpy of O or F is less than that of succeeding element. This is because when electron is added to O or F, the added electron goes to smaller n = 2 quantum level and suffers significant repulsion from the other electrons present in this level. For the n = 3 quantum level ( S or Cl) , the added electron occupies a larger region of space and the electron-electron repulsion is much less.

Some typical trends in Electron Gain Enthalpy

Halogens have high negative electron gain enthalpy. This is due to their strong tendency to gain an additional electron to change into s2p6 configuration. The electron gain enthalpy become less negative from Cl Br I , i.e., on moving down the group. However, the electron gain enthalpy of fluorine is unexpected less negative. It is probably due to small size of the atom(atomic radius 72 pm). The addition of extra electron produces high electron charge density in a relatively compact 2p sub-shell resulting in strong electron-electron repulsion. The repulsive forces between electrons imply low negative electron gain enthalpy value. Therefore the incoming electron is not accepted with the same ease as in the case of chlorine which the atomic size is bigger (atomic radius 99 pm) and the electron crowding is comparatively less. As a result, the negative electron gain enthalpy of fluorine (−328 k J mol−1) is less as compared to chlorine ( - 349 k J mol−1)

Electron Gain Enthalpies of Halogens

In the oxygen family (Group 16) the negative electron gain enthalpy of oxygen is also less than that of sulphur although it is expected to be more.

Apart from the size of the atom, the electronic configuration also influence their electron gain enthalpy values considerably.

The members of the noble gas family have highly symmetrical electronic configuration and have completely occupied orbitals. Therefore , there atoms have hardly any need or urge to take up extra electron. In case these are made to accept electrons under special conditions, enthalpy or energy will be needed and their electron gain enthalpies (ΔegH) will be positive. As we come down the group, the electron gain enthalpy decrease. However, the first element helium has exceptionally small electron gain enthalpy(48 k J mol−1) due to very small atomic size and high tendency to accept extra electrons.

Electron gain enthalpies of beryllium, magnesium, calcium and nitrogen are also positive. This is attributed to the extra stability of the fully completed s-orbitals in Be(2s2), Mg (3s2) and Ca(4s2 ) and of exactly half-filled 2p-orbitals in
N (2s2 Px1 Px1 Pz1 ). Thus, if an atom has fully filled or exactly half-filled 2p-orbitals, its electron gain enthalpy value will be positive.

Electronegativity

The term electronegativity has been defined differently by different investigators.

  1. Pauling's Scale : Pauling defined electronegativity as the power of atom in a molecule to attract electrons to itself.
  2. Mulliken's –Jaffe Scale : Mulliken suggested that the average of First Ionisation Energy (IE1 ) and Electron Affinity (EA) of an atom should be a measure of electronegativity of an atom. Accordingly, he proposed the following empirical correlation for calculating electronegativities of various element.
  3. Allred –Rochow scale : Allred and Rochow defined electronegativity as the force of attraction exerted by the nucleus of an atom on the valency electrons. Making use of the effective nuclear charge at the periphery of the atom, Zeff , they proposed the following empirical relation for calculating the electronegativity :
    where χ is the electronegativity and r is the covalent radius of the atom in angstrom units .

The electronegativity of any given element is not constant ; it varies depending on the element to which it is bound. Though it is not a measurable quatity, it does provide a means for predicting the nature of force that holds a pair of atoms together.

Factors affecting Magnitude of Electronegativity

Magnitude of electronegativity depends up on the following factors :

  • Size of the atom : Smaller the size of the atom, the greater is the attraction for bonding electrons. Thus, atoms with small size are more electronegative.
  • Nuclear charge : Electronegativity value increases with increase in nuclear charge. This is because the nucleus with higher nuclear charge attracts the shared pair of electrons more strongly towards itself.
  • Type of the ion : A cation attracts the electron pair strongly towards itself as compared to the atom from which it has been derived. This is due to the smaller size of the cation as compared to the neutral atom from which it has been derived. The cation has higher electronegativity than the parent atom. For example , the electronegativity of Li is 1.0 while that of Li+ is 2.5.

An anion attracts the electron pair less strongly as compared to its parent atom. This again is due to the fact an anion has larger size as compared to the parent atom from which it is derived. Thus, an anion has less electronegativity than the parent atom. For example, electronegativity of F atom is 4.0 while that of F is 0.8.

If an atom can form different types of cations, then the cation with higher positive charge has more value of electronegativity than the cation with lower positive charge. This is due to the fact that the cation with higher value of positive charge has smaller size and hence greater attraction for electrons. For example, electronegativity of Fe2+ is 1.83 and that of Fe3+ is 1.96 . Similarly electronegativity of Sn2+ is 1.81 and that of Sn4+ is 1.96.

Ionisation energy and electron affinity

The electronegativity of an element is the mean of ionisation energy and electron affinity. Thus elements having higher value of ionisation energy and electron affinity also have high values of electronegativity. For example, the elements of Group 1 (alkali metals) having lowest electron affinity and ionisation energy , have lowest electronegativity values. Similarly, elements of group 17 (halogens) which have high values of ionisaton energy and electron affinity have high electronegativity values.

Applications of Electronegaivities

Electronegativities have a very wide range of applications. Some of the important applications are given below :

  • Non-metallic elements have strong tendency to gain electrons. Therefore, electronegativity is directly related to that non-metallic properties of elements. Electronegativity is inversely related to the metallic properties of elements. Thus increase in electronegativities across a period is accompanied by an increase in non-metallic properties (or decrease in metallic properties) of elements . Similarly , the decrease in electronegativity down a group is accompanied by a decrease in non-metallic properties (or increase in metallic properties) of elements.
  • Calculation of partial ionic character of a covalent bond : The development of ionic character in a covalent bond between two atoms , say A and B , is due to the difference in the electronegativities of A and B. The greater the difference in the electronegativities, the greater would be the development of ionic character and consequently the higher would be the stability of the resulting bond.
  • Calculation of enthalpies of formation of compounds .
  • Calculation of bond length : If the two atoms A and B bonded together through a covalent bond differ in their electronegativities, then the covalent bond would acquire some ionic character, i.e., the bond acquires polarity. The greater the polarity, the shorter would be the length of the bond formed between A and B.
  • Explanation of bond angles : The lesser the electronegativity of the central atom in a polyatomic molecule, the lesser would be the bond angle.
  • Explanation of diagonal relationship : The electronegativity increases as we go from Li to Be (variation of electronegativity in a period) but it decreases as we move from Be to Mg (variation of electronegativity in a group). As a result of these two opposite changes(one along the period and other down the group), as we move diagonally , these two effects partly cancel each other and there is no marked change in electronegativity. This the reason why Li and Mg have similar chemical properties.
  • Type of bonds : The type of bond formed between two atoms would evidently, depend upon the differences in their electronegativities. If this difference is zero or very small, the bond formed would be covalent and and if the difference exceeds 2.5 , the bond formed would be electrovalent . If however , the electronegativity difference is less is less than 2.5 but otherwise quite appreciable, the bond formed will be polar covalent . The following table gives the relationship between the percentage ionic character and electronegativity difference.

Electronegativity difference

% Ionic character

0.1

0.5

0.2

1

0.6

10

1.1

25

1.7

50

2.3

75

3

90

Electronegativity Scale

The one which is the most widely used is Pauling scale. Pauling (1922) assigned arbitrarily a value of 4.0 to fluorine , the element considered to have the greatest ability to attract electrons. Approximate values of electronegativity of a few elements are given in the Table

Elecronegativities of the Main Group Elements

Group1

2

13

14

15

16

17

H

2.2

Li

1.0

Be

1.5

B

2.0

C

2.5

N

3.0

O

3.5

F

4.0

 

Na

0.9

Mg

1.3

Al

1.6

Si

1.9

P

2.2

S

2.6

Cl

3.1

K

0.8

Ca

1.0

Ga

1.6

Ge

1.6

As

2.2

Se

2.5

Br

2.9

Rb

0.8

Sr

1.0

In

1.8

Sn

1.8

Sb

2.0

Te

2.1

I

2.6

Cs

0.7

Ba

0.9

T l

1.6

Pb

1.8

Bi

2.0

Po

2.0

At

2.2

a) Variation along a Period.

Electronegativity increases along a period from left to right . Across the period, the combined effect of increased nuclear charge and decreasing atomic size results in an increasing tendency to attract the binding electrons in compounds.

b) Variation along a Group

In a given group, the electronegativity decreases with increase in atomic number. This is due to increased atomic size

Periodic Trends

All the Periodic Trends are summarised in the following Fig.

The periodic trends of elements in the periodic table

Anomalous properties of Second Period Elements

The first element of each of the group 1 (Lithium) and 2 (Beryllium) and groups 13 – 17 (boron to fluorine) differs in many respects from other members of their respective group. For example, lithium unlike other alkali metals and beryllium unlike other alkaline earth metals form compounds with pronounced covalent character ; the other members of these groups predominantly form ionic compounds. In fact the behaviour of lithium and beryllium is more similar with the second element of the following group i.e., magnesium and aluminium respectively. This sort of similarity is commonly referred to as diagonal relationship in the periodic properties.

Reasons for the anomalous behaviour

The anomalous behaviour is attributed to their small size, large (charge)/(radius) ratio and high electronegativity of the elements. In addition, the first member of the group has only four valence orbitals (2s and 2p) available for bonding, whereas the second member of the groups have nine valence orbitals (3s, 3p, 3d) . As a consequence of this, the maximum covalency of first member of each group is 4 (e.g. boron can only form [BF4]4−, whereas the other members of the groups can expand their valence shell to accommodate more than four pairs of electrons e.g. aluminium forms
[AlF6 ]3− . Further , the first member of p-block elements displays greater ability to form p π - p π multiple bonds to itself
( e.g. C=C , CC , N=N , NN) and to other second period elements (e.g. C=O, C=N, CN, N=O) compared to subsequent members of the same group.

 

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