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CHEM 2160 (C2OA)
CHEMISTRY OF THE MAIN GROUP ELEMENTS
Lecture 7
Tuesday, Nov. 3, 2009
r. rv n umar
(ROOM 306 C1; Ext 3261)
ma : rv n . umar s a.uw .e u.
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The Group 16 (Chalcogen) Elements
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Some covalent bond enthalpy terms (kJ mol-1)
, ,
and tellurium.
The most important allotrope of oxygen is O2
pale blue liquid or solid. In all phases, it isparamagnetic with a triplet ground state, i.e.
the two unpaired electrons have the same
.
-, .
catenation by sulfur is high and leads to the formation of both rings of varying sizes and chains. Allotropes of known structure
include cyclic S6, S7, S8, S9, S10, S11, S12, S18 and S20 (all with puckered rings) and fibrous sulfur (catena-S). In most of these,
the S-S bond distances are 2061 pm, indicative of single bond character; the S-S-S bond angles lie in the range 102108o.
The ring conformations of S6 (chair) and S8 (crown) are readily envisaged but other rings have more complicated
conformations. The structure ofS7 is noteworthy because of the wide range of S-S bond lengths (199218 pm) and angles(101.5107.58). The energies of inter-conversion between the cyclic forms are very small.
The most stable allotrope is orthorhombic sulfur (the -form and standard state of the element) and it occurs naturally as large
yellow crystals in volcanic areas. At 367.2 K, the -form transforms reversibly into monoclinic sulfur (-form). Both the - and
- - -3 -3 -. , .
of the rings is less efficient.
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Schematic representations of the structures of
some allotropes of sulfur: (a) S , (b) S , (c) Sand (d) catena-S
(the chain continues at each
end).
Nomenclature: The IUPAC namesfor H2O and H2S are oxane and
sulfane, respectively.2 = , , e, e
H2E hydrides are known for all the Group 16 elements, but become increasingly unstable on going down the
group, as shown by their enthalpies of formation (see table). This instability is paralleled by compounds such
as the EH3 series (E =N to Bi) and the HE series (E = F to I). The boiling points of the H2E compounds
revea e anoma ous y g o ng po n or wa er, ow ng o very s rong y rogen on ng. n mar e
contrast to water, H2S, H2Se and H2Te are highly toxic, foul-smelling gases.
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The hydronium or oxonium ion, H3O+ is well known. The sulfonium and
selenonium cations H3S+ and H3Se
+ can also be made (despite the lower
2 2 ,difficult) by using the very strong acid H+SbF6
- (generated from HF and SbF5)
H2S + HF + SbF5 H3S+SbF6-y rogen perox e, 2 2, is the most well-known peroxide. The
molecule adopts a gauche structure both in the gas phase and in the solidstate (see Figure), owing to repulsion of the lone pairs on the oxygen atoms.
Hydrogen peroxide undergoes a similar self-ionization to water, but to a
slightly greater extent. The equilibrium constant is 1.5 x 10-12, whereas for
water it is 10-14.
2H2O
2H
3O
2
+ + HO2
-
Hydrogen peroxide is stable towards decomposition into hydrogen and
oxygen, as shown by the large, negative enthalpy of formation (fHo = -187.8
kJ mol-1); however, it is unstable towards decomposition to water and oxygen
(see following equation).
2H2O2 2H2O + O2 H = -98.3 kJmol-1
The anthraquinolanthraquinone
process for hydrogen peroxide
n ustr a manu acture.
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In the laboratory, hydrogen peroxide can also be synthesized by the reaction of barium peroxide (BaO2) with
sulfuric acid and then removing the in soluble barium sulfate precipitate.
BaO2(s) + H2SO4(aq) H2O2(aq) + BaSO4(s)and M2E+2HCl H2E+2MCl (e.g., M2E=Na2S,Na2Se)
compoun s w c con a n peroxo groups are ox z ng agen s, an s s e ma n ea ure o e r
chemistry. No compounds are known which have greater than two oxygens in a chain terminated byhydrogens (such as HOOOH), but related fluorine-containing compounds such as FOOOOF and CF3OOCF,
are known.
num er o y r es are nown w c con a n one or more - -, - e- e- or - e- e- n ages. e onges
chains are found for sulfur, with compounds ofcomposition H2Sn being stable for up to 8 sulfuratoms, and
possibly more. This is not unexpected, since elemental sulfur itself has a strong tendency to form long chains
(catenation). n - + + - n- po ysu ane m x ure, n = -
A mixture of polysulfanes can be prepared by the addition of acid to a polysulfide solution. Compounds of the
type H-(E)n-H (E = S, Se or Te) tend to be unstable, giving the element and H2E.
- n-
2 + n-In the laboratory, H2S was historically prepared by a Kipps apparatus. The hydrolysis of calcium or barium
sulfides (see equation) produces purer H2S, but the gas is also commercially available in small cylinders.
FeS(s) + 2HCl H2S(g) + FeCl2(aq) (Kipps) a + 2 2 + a 2
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Hydrogen selenide may be prepared by following reaction, and a similar reaction can be used to make H2Te.
Al Se + 6H O 3H Se + 2Al(OH)The enthalpies of formation of H2S, H2Se and H2Te (see Table) indicate that the sulfide can be prepared by
direct combination of H2 and sulfur (boiling), and is more stable with respect to decomposition into its elements
than H2Se or H2Te.
Like H2O, the hydrides of the later elements in group 16 have bent structures but the angles of ~90o (see
Table) are significantly less than that in H2O (105o). This suggests that the E-H bonds (E = S, Se or Te) involvep character from the central atom (i.e. little or no contribution from the valence s orbital).
In aqueous solution, the hydrides behave as weak acids. The second ionization constant ofH S is ~10-19 and,
thus, metal sulfides are hydrolysed in aqueous solution. The only reason that many metal sulfides can be
isolated by the action of H2S on solutions of their salts is that the sulfides are extremely insoluble. For
example, a qualitative test for H2S is its reaction with aqueous lead acetate.H2S + Pb(O2CCH3)2 PbSblack t + 2CH3CO2H
Sulfides such as CuS, PbS, HgS, CdS, Bi2S3, As2S3, Sb2S3 and SnS have solubility products less than ~10-30
and can be precipitated by H2S in the presence of dilute HCl. The acid suppresses ionization of H2S, lowering
the concentration of S2 in solution. Sulfides such as ZnS, MnS, NiS and CoS with solubility products in the
range ~10-15 to 10-30 are precipitated only from neutral or alkaline solutions.
Protonation of H2S to [H3S]+ can be achieved using the superacid HF/SbF5. The salt [H3S][SbF6] is a white
crystalline solid which reacts with quartz glass; vibrational spectroscopic data for [H3S]+ are consistent with a
trigonal pyramidal structure like that of [H3O]+.
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o ysu anesPolysulfanes are compounds of the general type H2Sx where x 2. Sulfur dissolves in aqueous solutions
of group 1 or 2 metal sulfides (e.g., Na2S) to yield polysulfide salts, (e.g. Na2Sx). Acidification of such
solutions ives a mixture of ol sulfanes as a ellow oil, which can be fractionall distilled to ield H S x
= 26). An alternative method of synthesis, particularly useful for polysulfanes with x > 6, is by
condensation reaction (see reaction below)
2H S + S Cl H S + + 2HClThe structure ofH2S2 resembles that of H2O2 with an internal dihedral
angle of 91o in the gas phase. All polysulfanes are thermodynamically
unstable with respect to decomposition to H2S and S. They are usedor e prepara on o cyc o- n spec es.
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TheGroup17(Halogen)Elements
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The halogens, like all of the p-block elements, span a range of reactivities. This ranges from the mostelectronegative and most reactive element in the Periodic Table fluorine), to some of the least reactive, iodine
and astatine.
The -1 oxidation state occurs in the halide anions (e.g. CI- ). This oxidation state becomes increasingly
reducing on going down the group. Iodide, 1-, is a moderate reducing agent, while chloride shows few
reducing characteristics, except with very strong oxidizing agents. The trends in the stabilities of the maingroup halide compounds illustrate this, for example in Group 14. Pbl4 is a non-existent compound, owing to
the combination of an oxidizing metal centre, Pb(IV), and a reducing iodide anion. In contrast, PbCl4 and PbF4are more stable.
Positive oxidation states, +I, +3, +5 and +7, occur for chlorine, bromine and iodine, mainly in oxyanions and
interhalogen compounds. Compounds in the highest oxidation states generally contain the electronegative
elements oxygen and fluorine, e.g. IF7and IO4-.
Fluorine, with its small size and high electronegativity, coupled with the weakness of the F-F bond, means
that it is able to stabilize the very highest oxidation states of elements, e.g.AuF5, NiF4, PtF6.
Iodine (and to a lesser extent bromine) forms solvated cations (with iodine in the +1 oxidation state) such as
[I(pyridine)2]+, by reaction of I2, pyridine and AgNO3 in a non-aqueous solvent.
I2 + AgNO3 + 2.pyridine [I(pyridine)2]+NO3 + AgI
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In pure form all of the elemental halogens exist as diatomic molecules, X2. At room temperature, F2 and CI2are both gases (yellow and pale green, respectively), Br
2
is a deep red liquid, while I2
is a metallic purple
solid, showing the increasing intermolecular (van der Waals) forces for the heavier halogens.
For chlorine, bromine and iodine the bond dissociation energies decrease going down the group, owing to
poorer overlap between increasingly large atoms with diffuse orbitals. Fluorine has an anomalously low
on ssoc a on energy, ecause o ncrease repu s on o one pa rs on a acen uor ne a oms, ow ng o
their closer proximity in F2 compared to Cl2.
All of the elemental halogens are oxidizing agents, with the reactivity decreasing going down the group.
emen a uor ne s e mos reac ve o any e emen , an spon aneous y orms c em ca compoun s w
all the other elements except the lighternoble gases helium, neon and argon.
Reactions with fluorine must be carried out in special vessels, such as poly(tetrafluoroethene) (PTFE), inme a s suc as n c e w c orm a pass va ng ayer o e me a uor e, or n very ry g ass vesse s.
Chlorine and bromine are far less reactive than fluorine, but they still react with many elements directly,
while iodine is the least reactive of the four, and it often requires heating for reaction to proceed.
The synthesis of F2 cannot be carried out in aqueous media because F2 decomposes water, liberating
ozonized oxygen (i.e., O2 containing O3). Difluorine combines directly with all elements except O2, N2 and
e g er no e gases; reac ons en o e very v o en . any me a s are pass va e y e orma on o a
layer of non-volatile metal fluoride.
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The high reactivity of F2 arises partly from the low
bond dissociation energy and partly from thestrength of the bonds formed with other elements
It is extremely corrosive, being easily the most reactive element known. Inhalation causes irritation of the
2 . 2 , 2has an unpleasant smell and causes eye and respiratory irritation.
In the crystalline state, Cl2, Br2 or I2 molecules are arranged in layers as represented in following Figure.
The molecules Cl2 and Br2 have intramolecular distances which are the same as in the vapour (compare
cov, . 2 2layer are shorter than 2rv, suggesting some degree of interaction between the X2 molecules. The shortest
intermolecular XX distance between layers is significantly longer. In solid I2, the intramolecular II bond
distance is longer than in a gaseous molecule, and the lowering of the bond order (i.e. Decrease in.
that solid I2 possesses a metallic lustre and exhibits appreciable electrical conductivity at higher
temperatures; under very high pressure I2 becomes a metallic conductor.Part of the solid state structures of Cl2, Br2 and I2 in which molecules are arranged in stacked layers, and relevant
intramolecular and intermolecular distance data.
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Chan es in the ener levels of the MOs and the round state electronic
configurations of homonuclear diatomic molecules involving first-row p-block elements.
It is likely that the weakness of the FF bond is largely a consequence of repulsions between the
nonbonding electron pairs. The small size of the fluorine atom brings these pairs into close proximity when
F
F bonds are formed. Electrostatic repulsions between these pairs on neighboring atoms result in weakerbonding and an equilibrium bond distance significantly greater than would be expected in the absence of
such repulsions. In orbital terms, the small size of the fluorine atoms leads to poorer overlap in the formation
of bonding molecular orbitals and to improved overlap of antibonding * orbitals than would be expected by
extrapolation from the other halogens. For example, the covalent radius obtained for other compounds of
fluorine is 64 pm; an FF distance of 128 pm would therefore be expected in F2. However, the actual
distance is 143 pm. 13
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The observed colours of the halogens arise from an electronic transition from the highest occupied MO to
the lowest unoccupied MO. The HOMOLUMO energy gap decreases in the order F2 > Cl2 > Br2 > I2,
leading to a progressive shift in the absorption maximum from the near-UV to the red region of the visible
spectrum.
Dichlorine, dibromine and diiodine dissolve unchanged in many organic solvents (e.g. saturated
hydrocarbons, CCl4). However in, for example, ethers, ketones and pyridine, which contain donor atoms, Br2an 2 an 2 o a sma er ex en orm c arge rans er comp exes w e a ogen ac ng as e
acceptor orbital. In the extreme, complete transfer of charge could lead to heterolytic bond fission as in theformation of [Ipy2]
+. Solutions of I2 in donor solvents, such as pyridine, ethers or ketones, are brown or
yellow. Even benzene acts as a donor, forming charge transfer complexes with I2 and Br2; the colours of
ese so u ons are no cea y eren rom ose o 2 or r2 n cyc o exane a non- onor . ereas
amines, ketones and similar compounds donate electron density through a lone pair, benzene uses its
electrons; this is apparent in the relative orientations of the donor (benzene) and acceptor (Br2) molecules in
see Figure. That solutions of the charge transfer complexes are coloured means that they absorb in thev s e reg on o e spec rum nm , u e e ec ron c spec rum a so con a ns an n ense
absorption in the UV region (230330 nm) arising from an electronic transition from the solvent - X2 occupied
bonding MO to a vacant antibonding MO. This is the so called charge transfer band.TheevidencefortheweakeningoftheXXbondcomesfromvibrational spectroscopicdata,e.g.a
s t or
rom cm n2
to cm n6 6
.2.
1. ThereactionofPh3SbwithBr2 orI2 isanoxidativeadditionyieldingPh3SbX2.
2. Ph3AsBr2 isanAs(V)compound,whereasPh3AsI2,Me3AsI2 andMe3AsBr2 arechargetransfer
complexes ofthetypeshownin
. enatureo t epro ucts romreact onsare epen entont eso ventan t e group n
R3P.
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CH3CN.Br2
, , , 4 6 2. r2 2Ph3P.Br2
C6H6.Br215
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R3P + I2 R3PI4
P
iPr
I
iPr
iPr
PI
iPr
iPrI
II
368pm
iPr292pm
All the hydrogen halides, HX, are gases at 298K with sharp, acid
smells. Direct combination of H and X to form HX can be used H2 + X2 2HX (X = F, Cl, Br, I )synthetically only for the chloride and bromide. Even at low
temperatures, F2 reacts explosively with H2 in a radical chain
reaction. In the light-induced reaction of Cl2 and H2, the initiation stepis the homol tic cleava e of the ClCl bond to ive Cl radicals
which react with H2 to give H and HCl in one of a series of steps in
the radical chain; HCl can be formed in either a propagation or a
termination step. Reactions of H2 with Br2 or I2 occur only at higher
tem eratures and also involve the initial fission of the X molecule.
For Br2 (but not for I2) the mechanism is a radical chain.Hydrogen fluoride is prepared by treating suitable fluorides with
concentrated H2SO4 and analogous reactions are also a convenient
means of makin HCl.
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Analogous reactions with bromides and iodides result in partial
oxidation of HBr or HI to Br or I , and synthesis is thus byreaction 16.13 with PX3 prepared in situ. Halogens react with H2with the ease of reaction decreasing down group 17.
Hydrogen fluoride is an important reagent for the introduction of F
into organic and other compounds in the production of CFCs). It
differs from the other hydrogen halides in being a weak acid in
aqueous solution (pKa = 3.45). This is in part due to the high HFbond dissociation enthalpy. At high concentrations, the acid
strength increases owing to the stabilization of F by formation of
[HF2]-.
The formation of [HF2]- is also observed when HF reacts with group 1 metal fluorides; M[HF2] salts are stable
at room temperature. Analogous compounds are formed with HCl, HBr and HI only at low temperatures. .
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The melting and boiling points increase going down the group, with the exception of HF which has an
. ,where HF vapour mainly exists as a hexamer, (HF)6 up to 60
oC. In the solid state, HF exists in a polymeric
zigzag chain, (HF)n with strong F----HF hydrogen bonds.
.
HCI, HBr and HI are all very strong acids and are essentially fully dissociated in dilute solution; however, HF is
. , npolymeric chain such as H2F3
- and H3F4-, held together by the strong H----F hydrogen bonding that is so
important in this system.
Solvated Proton:The hydrogen can form the hydrogen ion only when its compounds are dissolved in media that solvate
protons. The solvation process thus provides the energy required for bond rupture, a necessary corollary of
this process is that proton H+
never exists in condensed phases, but occurs always as solvates (H3O+
, R2OH+
,etc).
H+(g) + xH2O = H+
(aq) H = -109 kJmol-1
, , .
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Factors affecting the acidity
The properties of aqueous solutions of acids are theproperties of the H3O+ ion, a solvated proton (hydrogen
ion) that is known as the hydronium ion in much of the older
c emca era ure u a so re erre o as e oxon um on.
Among the substances that react with water to produce
H3O+ are HCl, HNO3, H2SO4, HClO4, H3PO4, HC2H3O2, and
many others. Because water solutions of all of these3 , .
course there is a difference in degree because some are
strong acids while others are weak.
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AcidBase Equilibria. K as and pKas.
, . . .,
Theequilibriumconstantforthisreactionis:
BecauseH2Oisthesolvent,wedefineanewquantity,Ka:
The larger the value ofKa, the more likely an acid is to ionize, andhence the stronger that acid is.Because values for Ka vary by up to 60 orders of magnitude, we usually refer to them with a
logarithmic(common log, or log10) scale:
pKa = log(Ka) pKa(acetic acid) = log(1.8 x 105) = 4.7
The smaller the value of pKa, the stronger that acid is. Common mineral acids such as HCl have
pKa around 10. The pKa (H2O) = 15, while pKa (H3N) = 35, and pKa (CH4) 50. It is important to
see the relationship between the strength of an acid/base and the strength of its conjugate
base/acid. If an acid is very strong, then when it gives up its proton, the conjugate base will be very
unwilling to take that proton back, so it will be a weak conjugate base.
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strong acid gives weak conjugate base The pKa of acetic acid is 4.7, and the pKa of NH4+ is 10.
strong base gives weak conjugate acid
weak base gives strong conjugate acid
or examp e,CH3CO2H + NH3 CH3CO2
+ NH4+
There are several factors that affect acidity, and we will examine each in turn: electronegativity, size, charge,
inductive effects, hybridization, resonance, steric effects.
1. SizeWhen comparing two acids in which the protons to be given up are directly attached to two elements in the
same column of the eriodic table, the heavier element is more acidic than the lighter one, all other things
being equal. HI (10) > HCl (7) >> HF (3). This trend is contrary to electronegativity. The trend is due to the
increasingly pooroverlap between the tiny H(1s) orbital and the orbital of the acidic element as you go down
the periodic table. Electronegativity is not the sole determining factor in acidity!!! Size matters!
2. Electronegativity
When comparing two acids in which the protons to be given up are directly attached to two elements in the
same row of the periodictable, the more electronegative element is more acidic, all other things being equal,
because the more electronegative atom is better able to accept the electron pair in the conjugate base.
HF (3) > H2O (15) > NH3 (35) > CH4 (50); and H3O+ (0) > NH4
+ (10).
The phrase, all other things being equal is important: NH4+
> H2O, even though N is more electropositivethan O, because here the N has a formal +1 charge. Conversely, when comparing two bases that are in the
same row of the periodic table, the more electronegative atom makes the weaker base. So HO is a weaker
base than NH2, which is weaker than CH3
. Conversely, when comparing two bases that are in the same
column of the periodic table, the heavier atom makes the weaker base. So CH3S is a much weaker base
than HO, even though O is more electronegative.
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3. Charge
Given two acids with the same element, the one that has a positive charge is more acidic than the onethat is neutral. NH4+ (10) > NH3 (35), and H3O
+ (0) > H2O (15.7). When speaking about the acidity of
NH4+ and NH3, one needs to be careful. The pKa of NH4
+ is 10, and its conjugate base is NH3. The pKaof NH3 is 35, andits conjugate base is NH2
. If you are talking about the acidbase reactions of NH3,
you have to be very careful to distinguish whether NH3 is acting as a base or an acid. Likewise H2O, and
in fact any alcohol ROH and any amine RNH2 or R2NH.
4. ResonanceThe more good resonance structures can be written for the conjugate base of an acid, the acid is more
acidic. By far the most important resonance stabilizer is the carbonyl group. So CH3CO2H (4.7) >>
CH3CH2OH (16), and CH3CONH2 (17) >> CH3CH2NH2 (35); and CH3CHO (16.7) >> CH4 (48). In all
cases, a very good resonance structure can be drawn in which the negative charge on the conjugate
base is stabilized by resonance into the carbonyl. Two carbonyls have an even more pronounced effect
on acidity: so, CH3COCH2COCH3 (9) > CH3COCH3 (20). Different carbonyl groups have different
acidities, and the differences can be understood by looking at the stability of the C+O resonance
structure. The better this resonance structure, the less acidic the compound. So esters, in which this
structure is stabilized by resonance donation from O, are less acidic than ketones, which are not.
Another important resonance stabilizer is the Ph group. So PhOH (10) > Me2CHOH (17), and PhNH2
(27) > EtNH2 (35). Groups attached to the ortho and para positions of the Ph group can make an acideven more acidic. So 2,4,6-trinitrophenol (1) > PhOH (10). The nitro and CN groups are also important
resonance stabilizers; in fact, CH3NO2 (10) > CH3COCH3 (20), but CH3CN (28) < CH3COCH3 (20). The
effect of resonance stabilization in the nitro group is magnified by the inductive effect of three
electronegative atoms near the acidic C.
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5. Hybridization
For any given atom, acidity increases in the order sp3 < sp2 < sp hybridization. So HCCH (25) > H2C=CH2(37) > H3CCH3 (50). The reason is that when H+ is lost, the lone pair left behind in the conjugate base is
lower in energy when it is in an orbital with greater s character (because s- orbitals are lower in energy than
p- orbitals). Because hybrid orbitals are used to make bonds to H, the greater the s- character of that hybrid
orbital (50% in sp, 33% in sp2, 25% in sp3), the more acidic that H. Conversely, the more s character in the
orbital used to hold a lone pair, the less basic that atom is. So H3N (pKa of conjugate acid = 10) is morebasic than pyridine (pKa of conjugate acid = 5.2), which is more basic than MeCN (pKa of conjugate acid =
10), and H2O (pKa of conjugate acid = 0) is more basic than Me2C=O (pKa of conjugate acid = 7).
6. Inductive effects
Electronegative groups near the acidic atom make it a stronger acid. So CF3CO2H (0.5) > NCCH2CO2H
(2.5) > CH3CO2H (4.7), and CF3CH2OH (12.4) > CH3CH2OH (16), and CHCl3 (13.6) > CH4 (48), and
PhSO3H (6.5) > PHCH2OH (16). Note that because the acidic H is not directly bound to the
electronegative atoms, size has no influence here, and only electronegativity matters. Conversely, bases
with electronegative groups near the basic atom are less basic than those without. So C 6H5O (10) is less
basic than C6F5O (?).
7. Steric effects
Compounds with more steric bulk are less acidic than compounds with less steric bulk. So Me 3COH (19)
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, ,Resonance, then Hybridization, then Inductive effects, then Steric effects. Acronym: CRHIS. Follow the
scheme below.1. Identify the most acidic atom in each compound.
2. The acidic atoms are different.
a. If the atoms are in the same column, the compound with the heavier atom is more acidic, all other things
(CRHIS) being equal.b. If the atoms are in the same row, the compound with the rightmost atom is more acidic, all other things
(CRHIS) being equal.
. , ,
things (CRHIS) being equal.
d. If all other things (CRHIS) are not equal, compare pKas. Memorize the pKas of: HCl, CH3CO2H, NH4+, ROH,
CH3COCH3, HCCH, NH3, and CH4.
3. The acidic atoms are the same.
i. If the acidic atoms have different charges, the one with the more positive charge is more acidic.
ii. If the acidic atoms have the same charge, draw the conjugate bases. The compound whose conjugate base
has more resonance stabilization is more acidic.
iii. If the acidic atoms have the same charge and resonance stabilization but different hybridizations, the one with
> 2 > 3 .
iv. If the acidic atoms have the same charge, resonance stabilization, and hybridization, the compound with moreinductively electron withdrawing groups near the acidic atom is more acidic.
v. If the acidic atoms have the same charge, resonance stabilization, hybridization, and inductive effects, the
compound with less steric bulk near the acidic atom is more acidic.
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c y o nary y rogen ompoun s.AH(g)A-(g) + H+ (g)
(numerically the same as the proton affinity).
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Furtherreadings:
norgan c em stry y . . . .
InorganicChemistry by JamesE.HouseInorganicChemistry by P.A.Cox
Inorganic Chemistry byShriver&Atkins
ChemistryofElements by Greenwood
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