acidity, EN
Group 5A
Group 6A
Group 7A
CH4
NH3
H2O
HF
55
35
15.74
3.2
SiH4
PH3
H2S
HCl
35
27
7.1
-7
GeH4
AsH3
H2Se
HBr
25
23
3.8
-8
H2Te
HI
2.6
-9
EN
Group 4A
acidity, size
pKa’s of Binary Acids

Note that the arrows point in the direction of an increase.

Increasing EN for an acid (H-B) increases its acidity. Increasing size of H-B also increases acidity.

Acidity of binary acids increases left to right across each period as the EN increases.
Acidity of binary acids increases down each group as the size of the heteroatom increases greatly, even
though EN is increasing up the group.
Basicity trends of conjugate bases are opposite those of acidity trends of acids. See below.

basicity
pKb’s of Conjugate Bases of Binary Acids
Group 4A
Group 5A
Group 6A
Group 7A
OH-
F-
-21
-1.74
10.8
PH2-
HS-
Cl-
6.9
21
HSe-
Br
10.2
22
HTe-
I-
11.4
23
-
NH2
-41
SiH3-21
-13
CH3
GeH3
-11
-
-
AsH2
-9
-
basicity

-
pKa’s of Ternary (Oxy-) Acids
Ternary acids (oxyacids) contain hydrogen-oxygen-hereroatom links and conform to the general formula
HmXOn. These compounds are acidic when X is a nonmetal or metalloid but are basic when X is metallic.
With few exceptions (such as phosphorus oxyacids), H-atoms are bonded exclusively to O-atoms. Several
factors affect the strength of an oxyacid.
1. If there is more than one ionizable hydrogen, Ka1 > Ka2 > Ka3. The successive acidity constants differ by
ca. 5 powers of 10, i.e., 105, e.g., H3PO4, pKa1 = 2.1, pKa2 = 7.2, pKa3 = 12.4
2. Inductive and resonance effects are important. Acidity is greatest when the heteroatom is highly
electronegative and is increased by the presence of electron-withdrawing groups (e.g., -F or –CF3).
Thus the acidity increases left to right in the following series:
H3PO4 < H2SO4 < HClO4
HIO3 < HBrO3 < HClO3
CH3CO2H < FCH2CO2H < F2CHCO2H < CF3CO2H
3. Acidity increases with increasing number of oxygens. This is also an inductive (and resonance) effect.
The Ka’s increase successively in the following series by factors of ca.10 5, i.e., pka’s decrease by ca. 5:
HClO < HClO2 < HClO3 < HClO4

Electrophiles (E+) and Lewis acids are both electron pair acceptors (proton donors in the binary acids
shown below). They have similar periodic trends as seen in the following table. Note that the arrows
point in the direction of an increase.
acidity, E+

Group 4A
Group 5A
Group 6A
Group 7A
CH4
NH3
H2O
HF
55
35
15.74
3.2
SiH4
PH3
H2S
HCl
35
27
7.1
-7
GeH4
AsH3
H2Se
HBr
25
23
3.8
-8
H2Te
HI
2.6
-9
acidity, E+
pKa’s of Binary Acids
Nucleophiles (Nu:-) and Lewis bases are both electron pair donors, however their periodic trends are not
consistent (as seen in the following table). Note that the arrows point in the direction of an increase.
basicity, Nu:-
Group 4A
CH3
-
-41
SiH3

Group 5A
NH2
-
-21
-
PH2
-
Group 6A
Group 7A
OH-
F-
-1.74
10.8
HS-
Cl-
-21
-13
6.9
21
GeH3-
AsH2-
HSe-
Br
-11
-9
10.2
22
HTe-
I-
11.4
23
Nu: basicity
pKb’s of Conjugate Bases of Binary Acids
-
The size (polarizability) of nucleophiles has the largest effect on nucleophilicity. Larger, more polarizable
nucleophiles are more able to donate an electron pair (through distortion of their large valence electron
clouds).