Question

(a) Which of the following is the stronger BronstedLowry acid, \(\mathrm{HBrO}\) or \(\mathrm{HBr} ?(\mathrm{~b})\) Which is the stronger Brensted-Lowry base, \(\mathrm{F}^{-}\) or \(\mathrm{Cl}^{-}\) ? Briefly explain your choices.

Short answer

(a) HBrO is the stronger Bronsted-Lowry acid because its conjugate base (BrO\(^-\)) is more stable than the conjugate base of HBr (Br\(^-\)) due to resonance. (b) Cl\(^-\) is the stronger Bronsted-Lowry base because its conjugate acid (HCl) is less stable than the conjugate acid of F\(^-\) (HF) due to the greater bond strength between H and F in HF.

Step-by-step solution

Part (a): Comparing the acidic strength of HBrO and HBr

First, let's look at the conjugate bases formed after they lose a proton: - HBrO dissociates into BrO^- and H^+; - HBr dissociates into Br^- and H^+. Now, we will compare the stability of the resulting conjugate bases (BrO^- and Br^-, respectively). The BrO^- ion has an oxygen atom attached that has a lone pair of electrons available for resonance. Consequently, it delocalizes the negative charge over the O and Br atoms leading to greater stability. On the other hand, Br^- ion is a halide ion with a localized negative charge. Hence, BrO^- is more stable than Br^-. As the stability of the conjugate base increases, the acidity of the corresponding acid also increases. Therefore, HBrO is a stronger Bronsted-Lowry acid than HBr.

Part (b): Comparing the basic strength of F\(^-\) and Cl\(^-\)

First, let's look at the conjugate acids formed after they gain a proton: - F^- forms HF after gaining a proton; - Cl^- forms HCl after gaining a proton. Now, we will compare the stability of the resulting conjugate acids (HF and HCl, respectively). The bond strength between H and F in HF is stronger than the bond between H and Cl in HCl. This is because fluorine is more electronegative compared to chlorine, leading to a more polarized bond and greater bond strength. Hence, HF is more stable than HCl. As the stability of the conjugate acid increases, the basicity of the corresponding base decreases. Therefore, F^- is a weaker Bronsted-Lowry base than Cl^-. To summarize: (a) HBrO is the stronger Bronsted-Lowry acid. (b) Cl^- is the stronger Bronsted-Lowry base.

Key concepts

Acidic Strength

Acidic strength refers to how well a molecule can donate a proton (H\(^+\)) to a base. In the context of the Bronsted-Lowry theory, this is especially important. When comparing acidic strength, such as between HBrO and HBr, understanding the ability of the acid to dissociate and form its conjugate base is key.
For example, HBrO dissociates to form BrO\(^-\) and H\(^+\), while HBr forms Br\(^{-}\) and H\(^+\). It turns out that BrO\(^{-}\) is more stable because it can spread the negative charge over both bromine and oxygen through resonance. This stability makes HBrO a stronger acid than HBr.
A good rule of thumb is: the more stable the conjugate base, the stronger the acid. This is because a stable conjugate base suggests that the parent acid readily gives up its proton, showcasing its acidic nature.

Conjugate Base Stability

A conjugate base is what remains after an acid donates a proton. This stability is crucial in determining the acid's strength. The more stable a conjugate base, the better it supports the negative charge created by the loss of a proton.
In our example with HBrO and HBr, the BrO\(^{-}\) ion is more stable than the Br\(^{-}\) ion. The reason lies in resonance, where BrO\(^{-}\) can delocalize its charges.

• Delocalization of charge means the charge is spread out rather than being concentrated.
• Resonance involves sharing electrons across multiple atoms, creating several resonance structures.

These concepts explain why acids with more stable conjugate bases are stronger: they are more willing to lose a proton and transition into their stable conjugate base form.

Basic Strength

Basic strength is the ability of a molecule to accept a proton. In the Bronsted-Lowry framework, stronger bases are those that readily accept protons. Comparing F\(^{-}\) and Cl\(^{-}\), you can see how this works.
When you add a proton to F\(^{-}\), it forms HF. Likewise, adding a proton to Cl\(^{-}\) gives you HCl. However, the HF bond is exceptionally strong due to the high electronegativity of fluorine, which results in HF being quite stable and less likely to dissociate back into F\(^{-}\).
Thus, despite being very tough, HF's stability as a conjugate acid makes F\(^{-}\) a weaker base compared to Cl\(^{-}\). A simple guideline is: less stable conjugate acids correspond to stronger bases, as they suggest an increased tendency to gain protons.

Conjugate Acid Stability

Conjugate acid stability is key when examining the basicity of the corresponding bases. A conjugate acid forms when a base accepts a proton. The more stable this conjugate acid, the less powerful the base.
Take HF versus HCl. HF, formed from F\(^{-}\), is more stable than HCl, which forms from Cl\(^{-}\). The stability arises from a stronger bond between the hydrogen and the electronegative fluorine atom in HF.

• Increased bond strength means the conjugate acid (like HF) is less likely to dissociate.
• A stable conjugate acid implies a less reactive, weaker base.

This dynamic helps explain why the stability of conjugate acids is inversely proportional to the basic strength of their respective bases.