Сила кислотності у таблиці Менделєєва

Acidity Trends in the Periodic Table

We have previously seen that the strength of acids and bases is determined based on a molecule’s $K_{a}$ or $K_{b}$, respectively. Acids dissociate protons ($H^{+}$) and bases dissociate hydroxide ions ($OH^{-}$). The strongest acids and bases fully dissociate in water. But, within a molecule, there can be many different bonds to protons. Based on what atoms are around the protons we can determine the acidity of the protons and ultimately which is most likely to dissociate or be attacked by a basic molecule in a reaction.

In addition to $K_{a}$ values, the periodic table is an excellent tool to use to figure out the acidity of a molecule. The acidity of a molecule is based off its ability to donate a proton or accept an electron pair. When we think of the periodic table and its relation to acidity, what trends would be important?

Atomic size and electronegativity are the main trends that are important!

Out of these two trends, atomic size trumps electronegativity. Let’s also keep in mind strong acids have weak conjugate bases and completely dissociate in aqueous solutions. Strong bases have high electron density and a small size, to readily attack localizations of positive charge. Acidity increases across rows in the periodic table due to the increase in electronegativity and down periods due to the increase in atomic radius.

Electronegative atoms are responsible for pulling electron density around themselves from neighboring atoms and increasing the acidity of the neighboring atom.

This increased electron density around the electronegative atom leads to a partial negative charge on the electronegative atom and a partial positive on the atom whose electrons are being delocalized, refer to the above image to see the partial positive/negative charges. This delocalization of electrons increases the acidity of the protons. The greater the difference in electronegativity between an atom and its bonded proton, the greater the acidity.

Atomic size increases down the periodic tables’ periods. This increase in atomic size leads to an increase in atomic radius and ultimately a longer bond between the two atoms. For example, HI has a larger molecular bond than HF does. A longer bond is broken by less energy than a shorter bond, increasing the acidity of the molecule.

Note: $I^{-}$ is more electronegative than $F^{-}$, but due to the larger radius of $I^{-}$, HI is more acidic than HF.

Another great way of thinking about this is through the strength of the acids' conjugate base. In the Note, above, HF is less acidic than HI. The conjugate base of HF is $F^{-}$ and the conjugate base of HI is $I^{-}$. Because fluorine is more electronegative and smaller than iodine, it is the stronger base. Again, weak acids = strong conjugate base. This means that because $F^{-}$ is a stronger base, it has a weak conjugate acid, solidifying our knowledge on why an increase in molecular size increases acidity and weakens its ability to be a strong base.

$\color{#D73D75}{\text{Main Takeaways:}}$

$\color{#D73D75}{\text{1.Acidity increases across the periodic tables’ rows.}}$ $\color{#D73D75}{\text{2.Acidity increases down the periodic tables’ periods.}}$ $\color{#D73D75}{\text{3.Atomic size > electronegativity in importance to acidity.}}$

References: 1

(Video) Acidity Trends on the Periodic Table

By Chemistry Unleashed

In this video, Dr. Christiansen explains the acid-base periodic trends. He explains that acidity depends on the stability of the anion. He'll show you how stability changes across a period and down a group. You'll also discover why the stability changes. This is an excellent quick video for learning the periodic table's acid-base trend.