6.3.3: The acid-base behavior of binary element hydrides is determined primarily by the element's electronegativity and secondarily by the element-hydrogen bond strength
The compounds formed between the elements and hydrogen are called binary hydrides. All such compounds can in principle act as Brønsted acids in reactions with a suitably strong base. However, as the electronegativity decreases down a group and increases from left to right across the periodic table, the acidity of binary hydrides increases. In fact, on the left side of the periodic table the hydrides of extremely electropositive alkali and alkaline earth metals are not acidic but basic. They are perhaps best considered to be ionic salts of the hydride ion (\(\ce{H^{-}}\)). Consequently substances such as \(\ce{NaH}\) and \(\ce{CaH2}\) tend to act as Brønsted bases in their reactions.
\[\ce{NaH(s) + H2O(l) → Na^{+}(aq) + OH^{-}(aq) + H2(g)} \nonumber \]
\[\ce{CaH2(s) + 2H2O(l) → Ca^{2+}(aq) + 2OH^{-}(aq) + 2H2(g)} \nonumber \]
On the right side of the periodic table the binary hydrides of the nonmetals exhibit appreciable acidity.
\[\ce{HBr(aq) + H2O(l) → H^{+}(aq) + Br^{-}(aq)} \nonumber \]
For this reason the binary nonmetal hydrides are termed acidic hydrides. Nevertheless not all are equally acidic. The dilute aqueous acid ionization constants for these hydrides are given in Figure \(\PageIndex{1}\). As can be seen from the constants in Figure \(\PageIndex{2}\), the ability of the hydrides to transfer a hydrogen to water increases across a period and down a group.
These trends are largely due to changes in the electronegativity and size of the nonmetal atom:
- Foing across a period, the acid strength increases as there is an increase in electronegativity and the molecule gets more polar, with the hydrogen getting a larger partial positive charge. This makes it easier to heterolytically cleave the E-H bond to produce a stable anion. \[E-H → E:^- + H^+ \nonumber \]
- Going down a group, the acid strength increases because the bond strength decreases as a function of increasing size of the nonmetal, and this has a larger effect than the electronegativity. In fact \(\ce{HF}\) is a weak acid because it is so small that the hydrogen-fluorine bond is so strong that it is difficult to break. Remember, the weaker the bond, the stronger the acid strength. This is further illustrated in Table \(\PageIndex{1}\), where the weakest bond has produced the strongest acid.
| Relative Acid Strength | HF | << | HCl | < | HBr | < | HI |
|---|---|---|---|---|---|---|---|
| H–X Bond Energy (kJ/mol) | 570 | 432 | 366 | 298 | |||
| K a | 10 -3 | 10 7 | 10 9 | 10 10 |