Here, a charge is being ‘spread out’ (in other words, delocalized) by resonance, rather than simply by the size of the atom involved. Recall the fundamental idea that electrostatic charges, whether positive or negative, are more stable when they are ‘spread out’ than when they are confined to one atom. In the ethoxide ion, by contrast, the negative charge is ‘locked’ on the single oxygen – it has nowhere else to go. Chemists use the term ‘delocalization of charge’ to describe this situation. What this means is that the negative charge on the acetate ion is not located on one oxygen or the other: rather it is shared between the two. The two resonance forms for the conjugate base are equal in energy. For acetic acid, however, there is a key difference: a resonance contributor can be drawn in which the negative charge is localized on the second oxygen of the group. In both species, the negative charge on the conjugate base is held by an oxygen, so periodic trends cannot be invoked. How can they be so different in terms of acidity? We begin by considering the conjugate bases. In both compounds, the acidic proton is bonded to an oxygen atom. Comparing the other two to ethanoic acid, we see that phenol is very much weaker with a pK a of 10.00, and ethanol is so weak with a pK a of about 16 that it hardly counts as acidic at all! The pKa of ethanol is about 17, while the pKa of acetic acid is about 5: this is a 10 12-fold difference in the two acidity constants. Remember - the smaller the pKa, the stronger the acid. Three of the compounds we shall be looking at, together with their pK a values are: The smaller the number on this scale, the stronger the acid is. The strengths of weak acids are measured on the pK a scale.
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