The document discusses the concepts of acids and bases, focusing on Arrhenius, Brønsted-Lowry, and Lewis theories. It explains the distinctions between strong and weak acids and bases, their ionization in aqueous solutions, and introduces the concept of conjugate acid-base pairs. Additionally, it highlights the limitations of the Arrhenius definition and the broader applicability of the Brønsted-Lowry theory in various chemical reactions.

1. Acid Base concepts; Arrhenius,
Lowery-Bronsted, Lewis.
Soft and hard acid base
concept, Applications of SHAB

2. Arrheniusacids
□ The Arrhenius theory of acids and bases was originally proposed by the Swedish
chemist Svante Arrhenius in 1884. He suggested classifying certain compounds as
acids or bases based on what kind of ions formed when the compound was added to
water.
Citrus fruits—such as grapefruit—contain high amounts of
citric acid, a common organic acid

3. Arrhenius acid
An Arrhenius acid is any species that increases the concentration of
H+
ions—or protons—in aqueous solution. For example, let's consider the
dissociation reaction for hydrochloric acid, HCl, inwater:
When we make an aqueous solution of hydrochloric acid, dissociates into H+
ions. Since
this results in an increase in the concentration of H+
ions in solution, hydrochloric acid is an
Arrheniusacid.

4. The Arrhenius acid-base theory was proposed by Swedish Svante Arrhenius. It was the first
modern approach to acid-base concept. This theory is quite simple and useful.
According to Arrhenius theory, acids are the compound that increases the concentration
of H+
or proton in aqueous solution. The released H+
ion or proton is not free-floating proton,
itexistsin combined state with the water molecule and forms hydronium ion(H3O+
).
The common examples of Arrhenius acid includes HCl (hydrochloric acid), H2SO4
(sulphuric acid), HNO3 (nitric acid),etc.

5. Arrhenius acid formula Name
HClO3 Chloric acid
HNO3 Nitric acid
HClO4 Perchloric acid
H3PO4 Phosphoric acid
H2SO4 Sulphuric acid
H2SO3 Sulfurous acid
HCl Hydrochloric acid
CH3COOH Acetic acid
HBr Hydrobromic acid

6. When itisdissolved in water, then:
The acids like HNO3, HCl, etc. gives one proton on dissociation, called
monoproticacids.
Theacids like H2SO4, H3PO4, etc. which having more than one hydrogen atoms and
gives more than 1H+
ions on dissociation, called polyproticacids.
Itisnot necessary that polyprotic acids are stronger than monoprotic acids.

7. Arrheniusbases
□ Similarly, Arrhenius bases are compounds that increase the concentration of OH−
or
hydroxide ion in aqueous solution or having at least one OH−
ion in formula.
□ The common examples of Arrhenius base includes NaOH (sodium
hydroxide), KOH (potassium hydroxide), Ca(OH)2 (calcium hydroxide), Mg(OH)2
(magnesium hydroxide), NH4OH (ammonium hydroxide),etc.
□ When sodium hydroxide dissolved in water, it fully dissociates into ions Na+
and OH−
, this
dissociation increases the concentration of hydroxide ions in the solution.

8. Arrhenius base Name
NaOH Sodium hydroxide
NH4OH Ammonium hydroxide
KOH Potassium hydroxide
Mg(OH)2 Magnesium hydroxide
Ca(OH)2 Calcium hydroxide
Al(OH)3 Aluminum hydroxide

9. Neutralization reaction
When Arrhenius acid and Arrhenius base reacts, salt and water is formed as product, the
reaction isknown as neutralization reaction. Forexample:
□ The acids which are completely ionized in aqueous solution, is termed as strong acids
such as HCl, HNO3, H2SO4,etc.
□ Hydrochloric acid is a strong acid. When it dissociates into water, hydronium ion and
chloride ions are formed as product. Chloride ions are weak base, but its basicity does
not make the solution basic because acidity is overpowering the basicity of chloride
ions. The H+
ions combine with water molecule and form hydronium ion. In case of
strong acid, the concentration of hydronium ion formed is equal to the concentration
of the acid whereas in case of weak acids, the concentration of hydronium ions in
solution isalways lessthan the concentration of hydrogen ions.

10. □ Whereas the acids which are weakly ionized in aqueous solution, is termed as weak
acids such as acetic acid(CH3COOH).
□ In case of weak acids, the concentration of hydronium ion is always less than the
concentration ofacid.
□ Similarly, bases which are completely ionized in aqueous solution, are termed as strong
bases such as NaOH, KOH, etc. whereas the bases which are weakly ionized in
aqueous solution, is known as weak bases such as ammonium hydroxide (NH4OH),
calcium hydroxide (Ca(OH)2), etc.

11. Utilityof Arrhenius concept
□ Thistheory explains many phenomena like strength of acids and bases, salt hydrolysis
and neutralization.

12. Hydrogen ion (H+
) or hydronium ion(H3O+
)
□ When electron isremoving from hydrogen atom, hydrogen ion H+
isformed which isvery
reactive. But this H+
ion does not exist in aqueous solution. Since in aqueous medium, it
reacts with water molecule and forms hydronium ion (H3O+
). Water isa polar molecule;
it has the ability to attract the hydrogen ion (H+
). The water contains hydrogen and
oxygen in which oxygen (EN =3.5) ismore electronegative that pulls the electron density
towards it and causing the partial negative charge on the molecule. Due to partial
negative charge, it has ability to attract the positively charged hydrogen ion (H+
) and
form hydronium ion (H3O+
). Hydronium ions are more stable than hydrogen ions.
□ The hydronium ion is very important factor in chemical reaction that occurs in aqueous
solutions. Itisformed by the protonation of water.

13. LimitationsoftheArrhenius definition
□ The Arrhenius theory is limited in that it can only describe acid-base chemistry in
aqueous solutions. Similar reactions can also occur in non- aqueous solvents,
however, as well as between molecules in the gas phase. As a result, modern
chemists usually prefer the Brønsted-Lowry theory, which is useful in a broader range of
chemicalreactions.

14. S
ummary
□ An Arrhenius acid isany species that increases the concentration of H+
in aqueous
solution.
□ An Arrhenius base isany species that increases the concentration of OH-
in aqueous
solution.
□ In aqueous solution, H+
ions immediately react with water molecules to form
hydronium ions, H3O+
.
□ In an acid-base or neutralization reaction, an Arrhenius acid and base
usually react to form water and asalt.

15. Brønsted-Lowrytheoryofacids and bases
□ The Brønsted-Lowry theory describes acid-base interactions in terms of proton transfer
between chemical species. A Brønsted-Lowry acid is any species that can donate a
proton, H+
and a base is any species that can accept a proton. In terms of chemical
structure, this means that any Brønsted-Lowry acid must contain a hydrogen that can
dissociate as H+
. In order to accept a proton, a Brønsted-Lowry base must have at
least one lone pair of electrons to form a new bond with a proton.

16. □ Using the Brønsted-Lowry definition, an acid-base reaction is any reaction in which a
proton is transferred from an acid to a base. We can use the Brønsted-Lowry
definitions to discuss acid-base reactions in any solvent, as well as those that occur in
the gas phase. For example, consider the reaction of ammonia gas, NH3 with
hydrogen chloride gas, HCl, to form solid ammonium chloride, NH4Cl(s) :

17. This reaction can also be represented using the Lewis structures of the
reactants and products, as seenbelow:
In this reaction, HCl donates its proton—shown in blue—to NH3 Therefore, HCl is acting as
a Brønsted-Lowry acid. Since NH3 has a lone pair which it uses to accept a proton, NH3
isaBrønsted-Lowry base.
Note that according to the Arrhenius theory, the above reaction would not be an acid-
base reaction because neither species is forming H+ or OH− in water. However, the
chemistry involved-proton transfer from HCl to NH3 to form NH4Cl −very similar to what
would occur in the aqueous phase.

18. IdentifyingBrønsted-Lowryacids and bases
□ In the reaction between nitric acid and water, nitric acid, donates a
proton—shown in blue—to water, thereby acting as a Brønsted-Lowry acid.
□ Since water accepts the proton from nitric acid to form H3O+, water acts as a
Brønsted-Lowry base. This reaction highly favors the formation of products, so the
reaction arrow isdrawn only to theright.

19. Let'snow look at a reaction involving ammonia in water:
In this reaction, water is donating one of its protons to ammonia. After losing a proton,
water becomes hydroxide, OH−. Since water is a proton donor in this reaction, it is acting
as a Brønsted-Lowry acid. Ammonia accepts a proton from water to form an ammonium
ion, NH4+.Therefore, ammonia isacting as a Brønsted-Lowrybase.

20. Strongand weak acids: todissociate,or notto
dissociate?
□ A strong acid is a species that dissociates completely into its constituent ions in
aqueous solution. Nitric acid is an example of a strong acid. It dissociates completely
in water to form hydronium, H3O+ and nitrate, NO3− ions. After the reaction occurs,
there are no undissociated HNO3 molecules in solution.
□ By contrast, a weak acid does not dissociate completely into its constituent ions. An
example of a weak acid is acetic acid, CH3COOH which is present in vinegar.
Acetic acid dissociates partially in water to form hydronium and acetate ions,
CH3COO−:

21. □ Aqueous solutions of a strong acid, left, and a weak acid, right.
(a) Hydrochloric acid isa strong acid that fully dissociatesin water.
(b)Hydrofluoric acid isa weak acid that partially dissociates into protons and fluoride ions.

23. Strongand weak bases
□ A strong base isa base that ionizes completely in aqueous solution.An example of
a strong base issodium hydroxide, NaOH In water, sodium hydroxide dissociates
completely to give sodium ions and hydroxideions:
□ Thus,if we make a solution of sodium hydroxide in water, only Na+ and OH−
ions are present in our final solution. We don't expect any undissociated
NaOH.

24. Ammonia isa weak base, so itwill become partially ionizedin water:
Some of the ammonia molecules accept a proton from water to form ammonium ions
and hydroxide ions. A dynamic equilibrium results, in which ammonia molecules are
continually exchanging protons with water, and ammonium ions are continually donating
the protons back to hydroxide. The major species in solution is non-ionized ammonia, NH3,
because ammonia will only deprotonate water to a small extent.
Common strong bases include Group 1and Group2hydroxides.

25. Conjugate Acid-Base Pairs
□ Acids and bases exist as conjugate acid-base pairs. The term conjugate comes from
the Latin stems meaning "joined together" and refers to things that are joined,
particularly in pairs, such as Bronsted acids and bases.
□ Every time a Bronsted acid acts as an H+
-ion donor, it forms a conjugate base.
Imagine a generic acid, HA. When this acid donates an H+
ion to water, one product
of the reaction isthe A-
ion, which isa hydrogen-ion acceptor, or Bronstedbase.
HA + H2O H3O+
+
Acid +Base ⇌ conjugate acid +Conjugatebase
A-

26. The use of conjugate acid-base pairs allows us to make a very simple statement about
relative strengths of acids and bases. The stronger an acid, the weaker its conjugate
base, and, conversely, the stronger a base, the weaker its conjugate acid.

27. Amphoteric Compounds
□ A compound can be both a Bronsted acid and a Bronsted base. Water is the perfect
example of this behavior because it simultaneously acts as an acid and a base when
it forms the H3O+
and OH-
ions.

28. The Bronsted-Lowry acids and their
Conjugated Bases
□ Thestrength of the acid decreases as itdescends andthe strength of their
corresponding conjugate base increases.
Acids Conjugated base
Perchloric acid (HCIO4) Perchlorate ion (CIO4
–
)
Hydroiodic acid (HI) Iodide ion (I–
)
Hydrobromic acid (HBr) Bromide ion (Br–
)
Hydrochloric acid (HCl) Chloride ion (Cl–
)
Sulphuric acid (H2SO4) Hydrogen sulphate ion (HSO4
–
)
Nitric acid (HNO3) Nitrate ion (NO3
–
)
Hydronium ion (H3O+
) Water (H2O)
Hydrogen sulfate ion (HSO4
–
) Sulfate ion (SO4
2-
)
Nitrous acid (HNO2) Nitrite ion (NO2
–
)
Acetic acid (CH3COOH) Acetate ion (CH3COO–
)
Carbonic acid (H2CO3) Hydrogen carbonate ion (HCO3
–
)

29. Summary:
□ A Bronsted-Lowry acid isa substance which donates a proton or H+
ion to the other
compound and forms a conjugated base.
□ A Bronsted-Lowry base isa substance which accepts a proton or H+
ion from the other
compound and forms conjugatedacid.
□ Strong acids and bases ionize completely in an aqueous solution, whereas weak
acids andbases are partially ionized in aqueous solution.
□ Water molecule is amphoteric in nature, which means it can act as
Bronsted-Lowry acid as well as Bronsted-Lowrybase.

30. Lewis Acids and Bases
□ Lewis acids and bases are described by the Lewis theory of acid-base reactions as
electron-pair acceptors and electron pair donors respectively. Therefore, a Lewis base
can donate a pair of electrons to a Lewis acid to form a product containing a
coordinate covalent bond. This product is also referred to as a Lewis adduct. An
illustration detailing the reaction between a Lewis acid and base leading to the
formation of a coordinate covalent bond between them isgivenbelow.
□ Lewis acids and bases are named after the American chemist Gilbert Newton Lewis,
who also made invaluable contributions in the fields of thermodynamics and
photochemistry.

31. Lewis Acid
□ Lewis Acids are the chemical species which have empty orbitals and are able to
accept electron pairs from Lewis bases. This term was classically used to describe
chemical species with a trigonal planar structure and an empty p-orbital. An example
of such a Lewis acid would be BR3 (where R can be a halide or an organic
substituent).
□ Water and some other compounds are considered as both Lewis acids and bases
since they can accept and donate electron pairs based on the reaction.

32. Examples of LewisAcids
Some common examples of Lewis acids which can accept electron pairs include:
□ H+
ions (or protons) can be considered as Lewis acids along with onium ions like H3O+
.
□ The cations of d block elements which display high oxidation states can act as
electron pair acceptors. An example of such a cation isFe3+
.
□ Cations of metals such as Mg2+
and Li+
can form coordination compounds with water
acting as the ligand. These aquo complexes can accept electron pairs and behave
as Lewisacids.
□ Carbocations given by H3C+
and other trigonal planar species tend to accept
electron pairs.
□ The Pentahalides of the following group 15 elements can act as Lewis acids –
Antimony, Arsenic, and Phosphorus.
Apart from these chemical compounds listed above, any electron-deficient π system
can act as an acceptor of electron pairs – enones, for example.

33. Lewis Base
□ Atomic or molecular chemical species having a highly localized HOMO (The Highest
Occupied Molecular Orbital) act as Lewis bases. These chemical species have the
ability to donate an electron pair to a given Lewis acid in order to form an adduct, as
discussed earlier.
□ The most common Lewis bases are ammonia, alkyl amines, and other conventional
amines. Commonly, Lewis bases are anionic in nature and their base strength
generally depends on the pKa of the corresponding parent acid. Since Lewis bases
are electron-rich species that have the ability to donate electron-pairs, they can be
classified as nucleophiles. Similarly, Lewis acids can be classified as electrophiles
(since they behave as electron-pair acceptors).

34. Examples of LewisBases
Examples of Lewis bases which have an ability to donate an electron pair are listed below.
□ Pyridine and the derivatives of pyridine have the ability to act as electron pair donors. Thus, these
compounds can be classified as Lewis bases.
□ The compounds in which Oxygen, Sulphur, Selenium, and Tellurium (which belong to group 16 of
the Periodic Table) exhibit an oxidation state of -2 are generally Lewis bases. Examples of such
compounds include water andketones.
□ The simple anions which have an electron pair can also act as Lewis bases by donating these
electrons. Examples of such anions include H–
and F–
. Even some complex anions, such as the
sulfate anion (SO4
2-
) can donate pairs of electrons.
□ The π-systems which are rich in electrons (such as benzene, ethyne, and ethene) exhibit great
electron pair donatingcapabilities.
Weak Lewis acids have strong conjugate Lewis bases. Apart from this, many chemical species
having a lone pair of electrons such as CH3
–
and OH–
are identified as Lewis bases due to their
electron pair donatingcapabilities.

35. Chemical Reactions Between LewisAcids
and Bases
Reactionswiththe H+
ion
□ TheH+
ion acts as a Lewis acid and H2O acts as a Lewis base. Thereaction
between the water molecule and the proton yields a hydronium ion (H3O+
), as
illustrated below.
□ Here, the oxygen atom donates an electron pair to the proton, forming a
coordinate covalent bond in the process. The resulting Lewis acid has a+1charge
associated withit.

36. □ Another example of a reaction in which the H+
ion acts as a Lewis acid isits reaction
with ammonia (NH3) to form anammonium ion (NH4
+
).
□ In this reaction, the proton receives an electron pair from the nitrogen atom (belonging
to the ammonia molecule). The formation of a coordinate covalent bond between
the two resultsin the formation of a Lewis adduct (the ammonium cation).

37. Reaction Between Ag+
and Ammonia
□ In thisreaction, two Lewis bases form an adduct with one Lewis acid,as illustrated
below.
□ Here, ammonia acts as a Lewis base and the silver ion acts as a Lewis acid. Each
nitrogen atom donates an electron pair to Ag+
, resulting in two separate coordinate
covalent bonds. The adduct formed from the Lewis acid and base has the chemical
formula Ag(NH3)2
+
.

38. Reaction Between the Fluoride Ionand Boron
Trifluoride
□ Thisreaction features the formation of a coordinate bond between the fluoride anion
(F–
) and boron trifluoride(BF3).
□ Here, F–
acts as an electron pair donor whereas BF3 accepts the electron pair. The
reaction between the Lewis acid and base results in the formation of an adduct with
the chemical formula BF4
–
.

39. Applications of Lewis Acids and Bases
Some important applications of Lewis acids and bases are provided below.
□ Lewis acids play a vital role as a catalyst in the Friedel-Crafts reaction –
AlCl3 accepts a lone pair of electrons belonging to the chloride ion leading to the formation of
AlCl4
–
in the Friedel-Crafts alkylationprocess.
□ Thisalso leads to the formation of the highly electrophilic carbonium ion which acts as a strong
LewisAcid. The chemical reaction can be written as follows.
RCl+AlCl3 ⟶R+
+AlCl4
–
□ In the field of organic chemistry, Lewis acids are widely used to encourage many cationic or
pseudo-cationic chemical reactions.
□ Lewis bases have immense applications in the modification of the selectivity and the activity of
metallic catalysts. For the production of pharmaceuticals, asymmetric catalysis is an important
part of enantioselective synthesis. In order to enable asymmetric catalysis, chiral Lewis bases are
often used to confer chirality on catalysts.
□ Several Lewis bases have the ability to form many bonds with Lewis acids. These compounds are
also called ‘multidentate Lewis bases’ or ‘chelating agents’ and have a wide range of industrial
and agricultural applications.

40. Which Metals act as LewisAcids?
□ Metal ions such as Li+
and Mg2+
can accept pairs of electrons from a donating
species since they contain one or more empty orbitals. These ions tend to form
coordination compounds by accepting electron pairs from ligands.
□ Most metal ions exist in a coordinated structure with some ligands. For a Lewis base to
donate an electron pair to them, the metal ion must first dissociate from the ligand. The
Lewis adducts formed by these ions are generally complexes as well.

41. Some Examples of LewisBases.
A few examples of Lewis bases arelisted below.
□ Amines with the general formula R-NH3, such as methylamine.
□ Thefluoride ion (F–
)
□ Ammonia (NH3)
□ Water (H2O)
□ Acetone and many otherketones.
□ Compounds of the chalcogens (oxygen, sulphur, selenium, andtellurium) in which
they exhibit an oxidation state of -2 generally actas Lewis bases.

42. Isethyl acetate a Lewisbase?
□ Yes, ethyl acetate (or ethyl ethanoate) is a Lewis base since it has the ability to act as
an electron-pair donor. A Lewis base is a chemical compound that can donate a
pair of electrons to a suitable electron-pair acceptor (Lewis acid) to form a Lewis
adduct.
□ Thus, the definition, chemical behaviour, and the applications of Lewis acids and
bases are briefly discussed in thisarticle.

43. Hardand SoftAcids and Bases(HSAB)
Principle
HardandSoftAcids andBases(HSAB)Principleisa qualitative concept introduced by
Ralph Pearson to explain the stability of metal complexes and the mechanisms of their
reactions. However, itisalso possible to quantify this concept based on Klopman's FMO
analysis using interactions betweenHOMO and LUMO.

44. □ Lewis acids and bases can be classified by designating them as hard or soft. Acids
and bases are not strictly hard or soft, since many ions and compounds are classified
as intermediate. Thecategorizations are based on three factors:
(1) Charge density, or charge-to-sizeratio
(2) Polarizability
(3) Covalent vs.ionic nature of interactions.

45. According to HSAB principle, the Lewis acids and bases can be further divided into hard
or soft or borderlinetypes.
□ Hard Lewis acids are characterized by small ionic radii, high positive charge, strongly
solvated, empty orbitals in the valence shell and with high energy LUMOs.
□ SoftLewisacids are characterized by large ionic radii, low positive charge, completely
filled atomic orbitals and with low energyLUMOs.
□ HardLewisbasesare characterized by small ionic radii, strongly solvated, highly
electronegative, weakly polarizable and with high energyHOMOs.
□ S
oft Lewis basesare characterized by large ionic radii, intermediate
electronegativity, highly polarizable and with low energyHOMOs.
□ TheBorderlineLewis acids and bases have intermediate properties.

46. HSABPRINCIPLE
According to HSAB concept, hard acids prefer binding to the hard bases to give
ionic complexes, whereas the soft acids prefer binding to soft bases to give
covalent complexes. It is sometimes referred to as Hard-Soft Interaction Principle
(HSIP).

47. □ The large electronegativity differences between hard acids and hard bases give rise
to strong ionic interactions.
*The electronegativities of soft acids and soft bases are almost same and hence have
lessionic interactions. i.e., the interactions between them are more covalent.
*The interactions between hard acid - soft base or soft acid - hard base are mostly polar
covalent and tend to be more reactive or less stable. The polar covalent compounds
readily form either more ionic or more covalent compounds if they are allowed toreact.

48. HSAB&FrontierMolecularOrbitalTheory (FMO)
ANALYSIS
According to FMO analysis, the interactions between acids and bases are controlled by the relative
energies of the participating frontier molecular orbitals (FMO) i.e., HOMO and LUMO.
Greater the energy gap between the HOMO & LUMO, harder isthe species.
Quantitatively the absolute hardness of a species isdetermined by following equations.
or

49. Type of Acid/Base CHARACTERISTICS EXAMPLES
Hard acids
* Atomic centers of small ionic radii (<90 pm).
* High positive charge.
* Empty orbitals in their valence shells.
* Low electronegativity (0.7-1.6) and low electron affinity.
* Likely to be strongly solvated.
* High energy LUMO.
H+
, Li+
, Na+
, K+
, Be2+
, Mg2+
, Ca2+
, Sr2+
,Sn2+
Al3+
, Ga3+
, In3+
, Cr3+
, Co3+
, Fe3+
, Ir3+
, La3+
, Si4+
, Ti4+
,
Zr4+
, Th4+
,U4+
, VO2+
, UO2
2+
BeMe2, BF3, BCl3, B(OR)3, AlMe3
Soft acids
* Large radii (>90 pm).
* Low or partial positive charge.
* Completely filled orbitals in their valence shells.
* Intermediate electronegativities (1.9-2.5)
*Low energy LUMOs with a large magnitude of LUMO
coefficients.
Cu+
, Ag+
, Au+
, Hg+
, Cs+
, Tl+
, Hg2+
, Pd2+
, Cd2+
,
Pt2+
Metal atoms in zero oxidation states
BH3
Borderline acids intermediate properties
Fe2+
, Co2+
, Ni2+
, Cu2+
, Zn2+
, Pb2+
, B(CH3)3, SO2,
NO+
Hard bases
* Small radii (around 120pm) & highly solvated.
* electronegative atomic centres (3.0-4.0).
* Weakly polarizable.
* Difficult to be oxidized.
* High energy HOMO.
H2O, OH-
, F-
, Cl-
, CH3CO2-, PO4
3-
, SO4
2-
, CO3
2-
,
NO3
-
, ClO4
-
, ROH, RO-
, R2O, NH3, RNH2, N2H4
Soft bases
*Large atoms (>170 pm) with intermediate
electronegativity (2.5-3.0).
* High polarizability
* Easily undergo oxidation.
*Low energy HOMOs but large magnitude HOMO
coefficients.
S2-
, RSH, RS-
, R2S, I-
, CN-
, SCN-
, S2O3
-
, R3P, R3As,
(RO)3P, RNC, CO, C2H4, C6H6, R-
, H-
Borderline bases Aniline, pyridine, N3
-
, Br-
, NO2
-
, SO3
2-
, N2

51. APPLICATIONSOFHSABPRINCIPLE
There are numerous applications of the HSAB principle. It helps in understanding organic
reaction mechanisms, metal-ligand interactions in metal complexes, ore processing in
metallurgy, precipitations in qualitative analysis etc.
□ Inhydrogen bonding:
Thestrong hydrogen bond ispossible in cases of H2O, NH3, and HF,since the donor
atoms (F, O & N) are hard Lewis bases and their interactions with partially positively
charged H, which isa hard acid, are stronger.

52. Linkageofambidentateligandstometal atoms:
It is one of the important applications of the HSAB principle. The SCN-
ligand is an ambidentate
ligand and can be S-bound to metal (M-SCN) and referred to as thiocyanate or can be N-bound
to metal (M-NCS) and is referred to as isothiocyanate. The choice among S-bound or N-bound is
decided by soft or hard acid-base behavior. Sis a comparatively soft base than N atom. Hence soft
metal ions are S-bound while hard metal ions are N-bound.
1) SCN-
bonds through sulfuratom (soft base) when bonded to Pt2+
, a soft acid.
2) It bonds through nitrogen atom (a hard base) when linked to Cr3+
, a hard acid.
3)When Fe2+
reacts with SCN-
a bright red [Fe(SCN)]+
ion isformed, whereas Cr3+
forms [Cr(NCS)]2+
.
Reason: Fe2+
is a borderline acid and is S-bound. Whereas Cr3+
is hard acid and prefers to be N-
bound.
4)The molecule (CH3)2NCH2PF2 would bond to BF3through N whereas itwould bond to BH3through
P.
Reason: BF3is a hard acid and prefers to bind with N atom - a hard base. Whereas, BH3 is a soft
acid and preferentially bonded to soft base, P atom.

53. Symbioticeffect:
□ The hard-soft character of the metal ion is altered by the other groups attached. It is
referred to as a symbioticeffect.
□ For example, the isolated Co3+
is a hard acid and is expected to make the bond with
SCN-
ion through N atom as observed in[Co(NH3)5(NCS)]3-
.
□ However, when bound to five soft base ligands like CN-
ions, the hardness of cobalt
ion (Co3+
)is reduced. Thus [Co(CN)5]2-
behaves as a soft acid and prefers to bind with
SCN-
ion through Satom to form [Co(CN)5(SCN)]3-
.

54. Application ofHSABtopredictthe directionof
Inorganic reactions:
HSAB principle is used to predict the outcome of few of the reactions. We can predict
whether a reaction proceeds to the right or left based on soft or hard acid/base
interactions.
1)The reaction between AsF3 and PI3 is possible and proceeds to the right since As3+
is
softer than P3+
and I-
issofter than F-
.
Remember that both As3+
and P3+
are soft but relatively As3+
issofter due to larger size.
2)The reaction between MgS and BaO as shown below is possible since Mg2+
is harder
acid than Ba2+
and O2-
isharder base than S2-
.

55. □ P2F4 can be prepared by treating PF2I with mercury as shownbelow.
2PF2I +2Hg ------> Hg2I2 +P2F4
□ In thisreaction, itisiodine rather than fluorine that isremoved from PF2I.
□ Explanation: Hg2
2+
ion isa soft acid that prefers soft base I-
rather than hard base F-
.

56. Solubilityinwater:
□ The compound formed due to soft acid-soft base combination is more covalent and
less soluble in polar solvents like water. For example, Silver iodide, AgI is insoluble in
water as it has covalent nature since it is the combination of soft acid, Ag+
and soft
base, I-
.
□ On the other hand, Lithium iodide, LiI is the result of a combination of Li+
(hard acid)
and I-
(soft base). Thusitispolar covalent and thussoluble in water.

57. HardSoftinteractions- Typesofores:
□ We know that the hard metals prefer to bind with hard anions and thus they are
available as oxides or fluorides or carbonates or silicates in nature. Whereas, the soft
metals prefer to bind with soft anions and hence are found in nature as sulfides or
phosphides orselenides.
□ E.g.Aluminium ismostly found in nature as alumina, Al2O3 - an oxide ore, since Al3+
is
ahard metal which prefers to combine with hard oxide anion rather than the soft
sulfideion.
□ Silver& copper metals existas sulfide ores since both Ag+
and Cu2+
are soft metals.
□ Thef-block elements are found in nature as silicate minerals since the trivalent
lanthanides are actinides are hard acids and tend to bind with hard oxygen
bases as insilicates.

58. Precipitationreactions&Qualitative analysis:
□ The softer acids like Ag+
, Hg+
, Hg2+
etc., and borderline acids like Fe2+
, Ni2+
, Cu2+
, Zn2+
,
Pb2+
etc., can be precipitated as sulfides from their aqueous solutions since S2-
ion is a
softer base. Following table illustrates the separation of cations based on their
hardness orsoftness.