Question

(a) Which of the following is the stronger Bronsted-Lowry acid, HBrO or HBr? (b) Which is the stronger BronstedLowry base, \(\mathrm{F}^{-}\) or \(\mathrm{Cl}^{-} ?\)

Short answer

(a) HBr is a stronger Bronsted-Lowry acid than HBrO. (b) F⁻ is a stronger Bronsted-Lowry base than Cl⁻.

Step-by-step solution

(a) Compare acid strengths of HBrO and HBr

Consider the following reactions: HBrO ⇌ BrO⁻ + H⁺ (1) HBr ⇌ Br⁻ + H⁺ (2) To determine the stronger acid, we need to compare the stability of the conjugate bases (BrO⁻ and Br⁻) produced in the reactions.

(a) Stability of conjugate bases

BrO⁻ is a halogen-containing anion with a negative charge on the oxygen atom, whereas Br⁻ is a simple halide anion. Halogens are generally more electronegative than oxygen, so Br⁻ is significantly more stable due to electron distribution. Therefore, Br⁻ is more stable than BrO⁻.

(a) Stronger acid

As Br⁻ is a more stable conjugate base than BrO⁻, this implies that HBr is a stronger Bronsted-Lowry acid than HBrO.

(b) Compare base strengths of F⁻ and Cl⁻

To determine the stronger base between F⁻ and Cl⁻, we need to compare the stabilities of the conjugate acids produced from these bases: F⁻ + H⁺ ⇌ HF (3) Cl⁻ + H⁺ ⇌ HCl (4)

(b) Stability of conjugate acids

HF and HCl are both hydrogen halides. Since fluorine is more electronegative than chlorine, the H-F bond is more polar. This means that the HF molecule is less stable and has a weaker H-F bond compared to HCl.

(b) Stronger base

As HF is less stable as a conjugate acid than HCl, it implies that F⁻ is a stronger Bronsted-Lowry base than Cl⁻. In conclusion: (a) HBr is the stronger Bronsted-Lowry acid between HBrO and HBr. (b) F⁻ is the stronger Bronsted-Lowry base between F⁻ and Cl⁻.

Key concepts

Acid Strength

The strength of a Bronsted-Lowry acid is determined by its ability to donate a proton (H⁺). The stronger the acid, the more readily it will donate a proton. When comparing acids like HBrO and HBr, the focus should be on the stability of their conjugate bases. A strong acid will have a conjugate base that is stable, meaning it holds onto its negative charge effectively after the proton is donated.
For example, the conjugate base of HBr is Br⁻, while for HBrO it is BrO⁻. Since Br⁻ is more stable due to its simple structure and electron distribution, HBr is a stronger acid than HBrO.
Thus, understanding acid strength involves evaluating the ability of the conjugate base to handle or delocalize negative charge.

Conjugate Base Stability

Conjugate base stability plays a crucial role in determining the strength of an acid. A conjugate base is what remains after an acid donates a proton. The more stable this base is, the stronger the original acid.
Stability is often linked with the ability to spread out or delocalize the negative charge. In the case of Br⁻, the lack of extra oxygen atoms makes it a simpler anion, reducing electron repulsion and leading to higher stability. On the other hand, BrO⁻ has an extra oxygen atom which introduces more repulsion and less stability.
A stable conjugate base implies a strong parent acid because after donating a proton, the resulting anion isn't prone to revert back by re-acquiring a hydrogen ion.

Electronegativity Comparisons

Electronegativity is a key concept when assessing acid and base strength in terms of stability. It explains how strongly an atom attracts electrons, affecting both the acid and its conjugate base.
In acids like HBrO and HBr, the elements involved determine how effectively they stabilize their respective conjugate bases. Bromine is a less electronegative element compared to oxygen, thereby enabling Br⁻ to maintain a stable electron distribution.

• High electronegativity typically leads to stronger acids and weaker bases.
• With more electronegative elements like fluorine, the bonds are more polar, affecting acid and base strength differently.

For bases, more electronegative atoms can make the bond with hydrogen less stable, increasing the base's strength as it tends to hold onto electrons more.

Base Strength

Base strength in the Bronsted-Lowry framework is assessed by the base's ability to accept a proton effectively. The stronger the base, the more likely it is to capture a proton from its surroundings.
For \(F^−\) and \(Cl^−\), understanding which is the stronger base requires looking at their conjugate acids, HF and HCl. As HF forms a less stable molecule compared to HCl due to its high electronegativity, F⁻ is considered a stronger base than Cl⁻.
This inverse relationship shows that weaker conjugate acids make for stronger bases, as seen with F⁻ being a powerful base due to the instability of HF in comparison to HCl.

Conjugate Acid Stability

Conjugate acid stability is an important factor in determining the properties of a base. The conjugate acid is what forms when a base gains a proton. If the resultant conjugate acid is stable, the original base is considered weak because the gained proton isn't easily released.
HF and HCl serve as important comparisons. HF is less stable than HCl because of the strong polar bond between hydrogen and fluorine, making it difficult for the H-F bond within HF to dissociate.
Hence, since HF is less stable than HCl, its conjugate base \(\mathrm{F}^{-}\) is stronger than \(\mathrm{Cl}^{-}\). This showcases the balance between conjugate acid stability and base strength.