Covalent Bonding and Hard-Soft Acid Base Theory--Advanced Inorganic Chemistry

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					Covalent Bonding and Hard-Soft Acid Base Theory
Advanced Inorganic Chemistry

Covalent Bond Types
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Covalent bonds are characterized by a significant amount of electron sharing The atoms involved therefore have similar electronegativities, valence orbital energies and shapes. A key feature is that the orbitals involved must have positive overlap (I.e. the signs of the two orbitals match in the region of overlap; see figure 3.1 in text)

Three types of overlap
The net overlap is positive. This yields the bonding interaction that holds the two atoms together  The net overlap is zero. This is a nonbonding interaction  The net overlap is negative. This yields an antibonding interaction that forces the two atoms apart.
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Hybridization
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Originally grafted onto valence bond theory to account for tetrahedral structures, molecular orbital theory yields the same molecular orbitals Still commonly used in organic chemistry. Not used much in inorganic chemistry since it cannot account for many of the observed properties Hybridization involves mixing two or more sublevels (s and p commonly) to yield two new hybrid orbitals of different shape and orientation

Types of bonds - Sigma
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Sigma bonds -   no nodal plane through both atomic nuclei  strongest bonds with best overlap  electron density centered between the two nuclei  s+s, s+pz, pz+pz, etc.

Types of bond - Pi
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Pi bonds -   one nodal plane through both atomic nuclei  weaker bonds due to poorer overlap  bond strength falls off faster with increasing internuclear separation  electron density centers above and below internuclear axis  px+px, dxy+dxy, px+dxz

Types of bonds - Delta
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Delta bonds -   two nodal planes through both atomic nuclei  often occur in acetates and other derivatives of diatomic d-block elements 4+ 4+ 4+ 6+  Cr2 , Mo2 , W2 , Re2  dxy+dxy

Multiple bonds
double and triple bonds involve the addition of one or two -bonds to a -bond  quadruple bonds involve a -bond, two bonds and a -bond  multiple bonds are stronger than single bonds, but not two or three times stronger due to weaker overlap of the - and -bonds
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Hypervalent molecules
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in p-block atoms of 2nd period, the number of valence orbitals matches the maximum coordination number in 3rd and 4th periods, there are still only 4 valence orbitals, but as many as six at least partially covalent bonds are formed (PCl5 for example) these hypervalent molecules have more than an octet of electrons around the central atom

Two approaches
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Valence bond  six equivalent bonds require the hybridization of six atomic orbitals (d2sp3)  this is a hybrid involving two post-valent orbitals - the nd

Two approaches, con’t
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Molecular orbital  atomic orbitals can overlap with the atomic orbitals of more than one atom  generally, the nd postvalent orbitals of p-block central atoms are not used since they are too large and diffuse  sometimes they are included in compounds like SF6 but it is not clear how important they are

Hypervalency in d-block
there is no doubt that the np hypervalent orbitals are involved in bonding in d-block elements  particularly important later in the period  yields a total of nine valence orbitals  accounts for high coordination numbers
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Bond dissociation energies
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single bond energies tend to increase from left to right across a period, and up a group smaller atoms generally form stronger covalent bonds  shared electrons are closer to both atomic nuclei (and therefore more strongly attracted)  valence orbitals toward the bottom of the periodic table have more nodal spheres  valence orbitals of heavy atoms are more diffuse

Anomalies - too much of a good thing
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N-N, O-O, F-F bonds are much weaker than P-P, S-S and Cl-Cl bonds N, O, F are the smallest atoms so one would think that they would have the strongest bonds Yet consider H2N-NH2, H2O-OH2, F-F Each will have lone pairs in molecules. The smallness of the atoms causes the lone pairs to be very close, leading to significant electron-electron repulsion

The unique case of nitrogen
The N-N single bond is weak, so hydrazines and related molecules are easily oxidized (i. e. good reducing agents). Consider the following bond energies in kJ/mol:
B on d o rd er = C -C N -N O -O 1 3 47 2 93 1 59 2 6 11 6 15 4 18 3 8 37 8 91 9 46

Nitrogen - 1
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many nitrogen compounds are kinetically stable (have a high activation energy) The extraordinary strength of the nitrogennitrogen triple bond means that any reaction producing N2, once initiated, yields a heat of reaction sufficient to overcome the activation barrier and rapidly produce large amounts of hot expanding gas.

Nitrogen - 2
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The converse of this is that the production of reduced nitrogen compounds (ammonia, nitrates and nitrites) is very difficult. The stability of the nitrogen-nitrogen triple bond is unique, since both the N-N and N=N bonds are weak. The carbon-carbon triple bond is quite a bit more reactive, since it reacts readily with H2 , especially with a catalyst.

Explosives - 1
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These materials often exploit the nitrogen triple bond. Explosive are important to civil engineering (dams, highways, etc.), mining, explosive riviting, explosive metal forming, propellants, explosive bolts for aerospace, ice breaking and avalanche control, airbags, ejector seats and escape chutes.

Explosives - 2
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Explosives must be made of compound(s) that are only kinetically stable. The reactions must be highly exothermic, so that when only a few molecules react, the surrounding molecules can acquire the energy necessary to quickly react and send a detonation wave through the material. The hot reaction products must also be gaseous, so that the reaction does work on the surroundings through PDV expansion. This provides the destructive force, resulting in a nonreactive shock of ~103 m/s and pressures on the order of 10 GPa.

Explosives - 3
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Airports use neutron activation to screen for materials with high percentages of nitrogen. Commercial explosives are usually nitro compounds, which can oxidize the carbon to carbon oxides and the hydrogen to water vapor, increasing the exothermicity and gas yield of the reaction. The order of sensitivity to detonation is O-nitro>Nnitro>C-nitro. O-nitro compounds include dynamite, which is nitroglycerine absorbed onto kieselguhr. Cordite and gelignite are made by absorbing nitroglycerine onto nitrocellulose.

Explosives - 4
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Ammonium nitrate is stable to normal handling, and can even be gently heated to its melting point of 170 °C, where can decompose via the following reaction:

N H 4N O 3

N 2 O + 2H 2 O

Explosives - 5
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Above 250 °C or when shocked, the following reaction occurs:

N H 4N O 3

N 2 + H 2 O + 1/2 O 2

Explosives - 6
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Ammonium nitrate can be used as a blasting explosive alone or as a slurry in its own saturated aqueous solution Often, it is mixed with fuel oil to increase the explosive yield, as in the Oklahoma City bombing Since ammonium nitrate is the most widely used fertilizer, efforts are made to reduce its explosive potential Usually, it is mixed with calcium carbonate or ammonium hydrogen phosphates.

Explosives - 7
Initiators such as lead styphnate, lead azide and formerly mercury fulminate (Hg(ONC)2) are used to start the reaction.  These initiators are much less stable than the bulk explosives, but smaller amounts are needed to start the main reaction.
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Explosives - 8
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Black powder was known to the ancients, and is composed of 41% saltpeter (KNO3), 29.5% charcoal and 29.5% sulfur. A mixture of 87% KNO3 and 13 % charcoal also works. Modern black powder is 75 % saltpeter (KNO3), 15 % charcoal and 10 % sulfur. Generally speaking, any compound with a high nitrogen content is dangerous and potentially explosive.

Lewis Acid-Base Theory
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The presence of water actually diminishes the acidity of substances such as oxo acids and metal ions The water shields the species with hydration spheres In very concentrated solutions, the apparent concentration increases due to incomplete solvation For example, 19 M KOH behaves as if it were 10,000 M in hydrated hydroxide

Activity
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As we have considered before, a large number of water molecules are necessary to solvate cations such as Li+ As concentrations approach 0.1 M, there are insufficient solvent molecules available to shield the various ions in solution from each other. Analytical and physical chemists have to account for the deviation of concentrated solutions from the Law of Chemical Equilibrium.

Activity
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The concept of activity is used instead of concentration, where :

 

apparent

concentrat ion

actual concentrat ion

Activity
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The activity coefficient of the dissolved ion increases as the concentration of the ion increases due to the lack of sufficient solvent to properly shield the ion. Gas phase acidities and basicities are helpful in understanding this effect In the end, we must remember that the actual (unsolvated, gas-phase) behavior of an ion is not dependent on the presence of water, but is an intrinsic property of the ions

Lewis Acid-Base theory
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The growth of chemical investigations in nonaqueous and mixed-solvent media has increased the importance of considering acid-base behavior in the absence of water Lewis expanded the definitions of acids and bases to include non-protic, non-hydroxide systems  Lewis Acid - electron pair acceptor; electrophile  Lewis Base - electron pair donor; nucleophile

Lewis Acid-Base Theory
Lewis acids and Lewis bases react to form a coordinate covalent bond, where both electrons are supplied by the Lewis base  This species is referred to as an acid-base adduct, coordination complex or complex ion.
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Identifying Lewis Acids and Bases
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To determine Lewis acid-base character:  Draw the Lewis dot structure  Assign formal charge/oxidation state  Consider structures below

Classes of Lewis bases (Ligands)
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Monodentate Ligands - A ligand that donates only one electron pair to a single metal  One of the best examples is NH3. Bridging Ligands - A ligand that donates one or more electron pairs to two or more metals  Halides and hydroxide are good examples, since each possesses two or more electron pairs on the donor atom  As we considered before, one of the steps to forming insoluble hydroxide and oxo compounds is the bridging of two metals in the course of hydroysis.

Classes of Lewis bases – con’t
Ambidentate Ligands - Ligands that possess two or more donor atoms can act as a monodentate ligand through either donor atom, or bridge two metals.  The pseudohalides such as cyanide, azide and thiocyanate are good examples.
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Classes of Lewis bases – con’t
Chelating Ligands - Ligands with two or more donor atoms that can simultaneously interact with a single metal atom.  They must be nonlinear and form bonds with reasonable angles (90° or 109.5°).  Generally, the most stable chelating ligands form five- or six-membered rings
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Chelating ligands
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ethylenediamine (on the left) yields a fivemembered ring acetylacetonate (on the right) yields a sixmembered ring

N Ni N Ni

O

O

Denticity
The number of potential donor interactions is called the denticity  ligands are classified as bidentate chelating ligands (as for ethylenediamine), hexadentate chelating ligands (as for ethylenediaminetetraacetic acid; EDTA) and so forth.
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Macrocyclic ligands
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Macrocyclic Ligands - A special class of generally tetradentate chelating ligands, where the four donor atoms are arranged in a circular ring array, as in the porphyrins and anes.

14

Thermodynamics of Lewis Acid-Base Interactions
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Drago's E and C Parameters  The energies of Lewis acid-base reactions were investigated in the gas phase or in very nonpolar solvents, which approximate gas phase behaviors.  The Drago-Wayland Equation was formulated to predict such behavior: DH (kJ/mol) = -4.184 (CACB + EAEB) CA,CB, EA,EB are empirical parameters

E and C parameters
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While the Drago-Wayand equation is accurate only for uncharged species, it is instructive to consider the effects of the parameters. The E parameter is a measure of the strengths of the electrostatic contribution to the bonding. The C parameter is a measure of the strengths of the covalent contribution to the bonding. The E and C parameters allow us to compare the abilities of species such as I2 and SbCl5 to form covalent bonds.

E and C parameters – con’t
The CA values for I2 and SbCl5 are 1.00 and 5.13, respectively  One can say that SbCl5 is "better" at covalent bonding than I2  The key point is that it is impossible to characterize the Lewis acid-base interaction by a single parameter such as strength.
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E and C parameters – con’t
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When we considered the relative acidities of metal ions, we actually referenced them to the common base H 2O. Therefore, only one parameter was needed (we weren't comparing the acidities to a range of bases as we must in the Lewis situation). Lewis behavior always involves mixes of electrostatic/ionic and covalent contributions to the bonding, so at least one parameter is needed for each species. In charged species, a third parameter is needed to assess the degree to which the extent to which the negative charge of the anion is transferred to the cation in the adduct (charge polarization).

Hard-Soft Acid-Base Concepts
Hard Lewis acids tend to combine with hard Lewis bases, and soft Lewis acids tend to combine with soft Lewis bases  The important properties to consider in the classification of Lewis acids and bases are electronegativity, size and charge; the key features of many of our deliberations.
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Soft Acids
heavy metal ions  insoluble chlorides  chemically inert  fairly electronegative (1.9 - 2.54)  large size (> 90 pm)  low charge (+1 or +2)
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Hard Acids
most metals  electronegativity between 0.7 and 1.6  small (<90 pm)  often highly charged (3+ or higher)
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Soft Bases
C, P, As, S, Se, Te, Br, I donor atoms  electronegativity between 2.1 and 2.96  radius > 170 pm
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Hard Bases
O, F donor atoms  very high electronegativity  r ~ 120 pm  typical examples are sulfate, carbonate, silicate, acetate, alcohols, ketones
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Principle of Maximum Hardness
Reacting molecules will arrange their electrons so as to be as hard as possible (the electrons are more tightly held to the atoms and are therefore less polarizable).  Generally, hard molecules are those with large HOMO-LUMO gaps, and soft molecules are those with small HOMOLUMO gaps.
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Relative Softness
The relative softness of soft acids can be estimated from the distance of the acid from Au, the softest atoms.  For 1st row metals, the 2+ state is borderline,the 3+ state is clearly hard and the 1+ state (only observed for traditional coordination complexes as Cu(I)) is soft.  Acids get softer as they get heavier.
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Relative Softness
The softest bases are those with the lowest electronegativity and lie along the metalnonmetal border.  Donor atom softness is also influenced by the substituent atoms in bases like phosphines and arsine.
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Example of the use of the HSAB Principle
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Predicting to which side and equilibrium lies:

Cation Nb5+ Hg2+

Type Why hard soft High electronegativity and charge Low electronegativity and charge; closer to Au

Example of the use of the HSAB Principle
Anions Type Why S2O2soft Lower electronegativity, large

Hard High electronegativity, small

Therefore, products are favored since they match hard with hard and soft with soft.

Electronegativity Revisited
Hard Acid:Hard Base reactions usually have small enthalpy changes and large entropy changes

This suggests significant ion-solvent interactions are involved and the principle molecular interaction is ionic.

Electronegativity Revisited
Soft Acid:Soft Base reactions usually have large enthalpy changes and small entropy changes, consistent with the formation of a covalent bond:

Electronegativity within a molecule
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Pauling examined the reactions of elements and halides He separated the bonding into ionic and covalent contributions The ionic contribution is proportional to Z2/r Overall, the calculation of element (E)-halide (X) interactions can be formulated as:

Electronegativity within a molecule
One application of this is to estimate the electronegativity of an atom within a molecule.  For example, I in IF7, IF5, IF3, IF.  The electronegativity of F is 3.98  the I-I bond energy is 149 kJ/mol  the F-F bond energy is 155 kJ/mol  The average bond energy is 152 kJ/mol
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Electronegativity within a molecule
Oxidation Number p
IF7 IF5 VII V 3.07 2.88

IF3 IF

II I

2.87 2.84

Electronegativity within a molecule
Generally, electron donating substituents (low electronegativities) lower the electronegativity of an atom  Electron withdrawing substituents (high electronegativities) raise the electronegativity of an atom  The influence of the oxidation state on electronegativity is small.
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Strength and Softness
We need to characterize acid-base behavior with two parameters (at least)  We have addressed strength with Z2/r and other considerations  The general strength rule is that the strongest acid reacts with the strongest base, and the weakest acid reacts with the weakest base.
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Conclusions
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Remember that solvated ions (H+) have as their base the solvent. Think about solvated species as an acid-base pair that remains in solution, such as the OH2:OH- pair above. Precipitated ions have counterions other than solvent. This considerations are only for species that are sigmabonded To some extent, strength and softness are related through size, charge and electronegativity properties. Changes in substituents may change other properties that will then dominate the reaction, superceding acid-base considerations.


				
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